Full text of "Elementary Inorganic Chemistry"

See other formats

Google

This is a digital copy of a book that was preserved for generations on Hbrary shelves before it was carefully scanned by Google as part of a project

to make the world's books discoverable online.

It has survived long enough for the copyright to expire and the book to enter the public domain. A public domain book is one that was never subject

to copyright or whose legal copyright term has expired. Whether a book is in the public domain may vary country to country. Public domain books

are our gateways to the past, representing a wealth of history, culture and knowledge that's often difficult to discover.

Marks, notations and other maiginalia present in the original volume will appear in this file - a reminder of this book's long journey from the

publisher to a library and finally to you.

Usage guidelines

Google is proud to partner with libraries to digitize public domain materials and make them widely accessible. Public domain books belong to the
public and we are merely their custodians. Nevertheless, this work is expensive, so in order to keep providing this resource, we liave taken steps to
prevent abuse by commercial parties, including placing technical restrictions on automated querying.
We also ask that you:

+ Make non-commercial use of the files We designed Google Book Search for use by individuals, and we request that you use these files for
personal, non-commercial purposes.

+ Refrain fivm automated querying Do not send automated queries of any sort to Google's system: If you are conducting research on machine
translation, optical character recognition or other areas where access to a large amount of text is helpful, please contact us. We encourage the
use of public domain materials for these purposes and may be able to help.

+ Maintain attributionTht GoogXt "watermark" you see on each file is essential for informing people about this project and helping them find
additional materials through Google Book Search. Please do not remove it.

+ Keep it legal Whatever your use, remember that you are responsible for ensuring that what you are doing is legal. Do not assume that just
because we believe a book is in the public domain for users in the United States, that the work is also in the public domain for users in other
countries. Whether a book is still in copyright varies from country to country, and we can't offer guidance on whether any specific use of
any specific book is allowed. Please do not assume that a book's appearance in Google Book Search means it can be used in any manner
anywhere in the world. Copyright infringement liabili^ can be quite severe.

About Google Book Search

Google's mission is to organize the world's information and to make it universally accessible and useful. Google Book Search helps readers
discover the world's books while helping authors and publishers reach new audiences. You can search through the full text of this book on the web

at |http : //books . google . com/|

F-Juj-tT -J-oT-l-^ . a I . tft. S^

HARVARD UNIVERSITY

LIBRARY OF THE

Department of Education

COLLECTION OF TEXT-BOOKS

Contributed by the Publisbets

TRANSFERRED

LLEGE

3 2044 102 874 393

Bell's Science Series

Edited by
Percy Groom, M.A., D.Sc, and G. M. Minchin, M.A., F.R.S.

ELEMENTARY .
INORGANIC CHEMISTRY

BELL'S SCIENCE SERIES

ELEMENTARY BOTANY. By Percy Groom, M.A., D.Sc,
F. L. S ., Lecturer on Botany at Cooper's Hill, sometime Examiner
in Botany to the University of Oxford. Third Edition. With
275 Illustrations. Crown 8vo, 3s. 6d.

THE STUDENT'S DYNAMICS. Comprising Statics and
Kinetics. By G. M. Minchin, M.A., F.R.S., Professor of
Applied Mathematics at Cooper's Hill. 3s. 6d.

ELEMENTARY INORGANIC CHEMISTRY. By James
Walker, D.Sc, Ph.D., F.R.S., Professor of Chemistry in
University College, Dundee. 3s. 6d.

AN INTRODUCTION TO THE STUDY OF THE
COMPARATIVE ANATOMY OF ANIMALS. By G. C.
- Bourne,' M.A., Fellow and Tutor of New College, Oxford.
With numerous Illustrations.

Vol. I. Animal Organisation. The Protozoa and Coelen-
terata. 4s. 6d.

Vol. II. The Coelomata. 4s. 6d. [In the Press.

ELEMENTARY GENERAL SCIENCE. By D. E. Jones,
D.Sc, formerly Professor of Physics in the University College
of Wales, Aberystwith, and D. S. M'Nair, Ph.D., B.Sc

[/« the Press.

PHYSIOGRAPHY. By H. N. Dickson, F.R.S.E., F.R.Met.
Soc. , F. R. G. S. \In the Press.

ELECTRICITY AND MAGNETISM. By Oliver J. Lodge,
D.Sc, F.R.S.,LL.D., M.I.E.E., Principal of the University
of Birmingham. [/« Preparation.

ENTOMOLOGY. By L. C. Miall, F.R.S., Professor of
Biology in Yorkshire College, Leeds. [/« Preparation.

LIGHT. By A. E. Tutton, B.Sc, F.R.S. [/« Preparation.

LONDON: GEORGE BELL & SONS

YORK STREET, COVENT GARDEN

ELEMENTARY
INORGANIC CHEMISTRY

LONDON

GEORGE BELL & SONS

1901

Aaa*

-\ 1 1-^ X'^*\ '

' Hap/ard University,
Dept. of.Edowdtlon Libraryi
Gift of the Publisher!.

1:^,.[ > L.. -.D TO
HARVARD CULLtbE LIBRARY
Ar.i 16 1921

PREFACE

Of late years the teaching of Chemistry in schools has
tended ta assume more and more the character of " research
work," each pupil "discovering" his own facts (with a little
guidance), and also interpreting them in their mutual
connections. This method is admittedly of great value in
cultivating the -pupil's powers of accurate observation and
manipulation; but from the point of view of the student
who proceeds to a College or University, there to continue
the subject as part of a course in Science or Medicine, it
leaves something to be desired. In even an elementary
University course, the fitudent is suddenly confronted by a
bewildering multitude of de,tails/ < which he must learn from
his lecture-notes or a text-book, for now the limited time
at his disposal for practical work makes it impossible for
him to acquire his material on the leisurely school lines to
which he has become accustomed.

This little volume has been written to help in bridging
the gap here apparent. In it I have emphasised general
principles, which shall enable the student to extend his first-
hand laboratory knowledge in various directions, and to
connect and correlate apparently isolated facts. No instruc-
tions for practical work are given, as any teacher can easily
select from the many excellent manuals at his disposal those
experimehts which best typify the principles referred to.

Although I have written the book throughout from the
point of view of modern theory (so that the student on
pursuing his studies shall have as little as possible to unlearn),
I have studiously placed theoretical matters in the back-
ground, for the reason that in my own experience the beginner
does not appreciate their proper significance. The school
is, in my opinion, no place for " chemical philosophy " : that
should only come at the end of an elementary University
course, when the student has a good grip of the facts and
principles, i.e, the realities, of Chemistry.

J. W.

November 1901.

CONTENTS

Chapter

PAGE

Examples of Chemical Action .

II.

Conditions of Chemical Action .

III.

Solutions and Solubility

IV.

Symbols — Formulae — Equations

Combustion .....

VI.

Flame ......

VII.

Neutralisation .....

VIII.

The Common Acids and Bases

IX.

Oxides ......

Formation and Decomposition of Salts ,

XI.

Positive and Negative Radicals

XII.

Double Decomposition ....

XIII.

lonisation and Displacement of Radicals

XIV.

Electrolysis .....

XV.

Examples of Chemical Transformation

XVI.

Oxidation and Reduction

XVII.

The General Laws for Gases .

108

XVIII.

Gaseous Mixtures ....

117

XIX.

The Atmosphere . . . . .

122

XX.

Oxygen. . .

127

XXI.

Water .......

134

XXII.

Carbon .......

143

XXIII.

Nitrogen

154

XXIV.

Hydrc^en ......

165

XXV.

Chlorine . . . .

170

XXVI.

Bromine and Iodine . . . . .

182

XXVII.

Sulphur. ......

191

XXVIII.

Phosphorus ......

206

XXIX.

Silver — Copper — Mercury . . . .

220

XXX.

Lead — Tin ......

230

XXXI.

Zinc — Aluminium . . . . .

236

XXXII.

Iron .......

242

XXXIII.

Calcium — Barium . . . . .

248

XXXIV.

Sodium — Potassium — Ammonium

252

Index

,♦ » • » t » • •

261

\i INORGANIC CHEMISTRY

CHAPTER I

EXAMPLES OF CHEMICAL ACTION

If a piece of limestone is heated to bright redness, it under-
goes a complete change in its properties : it, in fact, ceases to
be limestone, and becomes what we call quicklime. The
change here produced by the action of heat is a chemical
change, and one of a comparatively simple character, so that
we may devote some time to its study.

When we say that a substance has been changed into some-
thing else, we mean that the properties of the material before
and after the change are recognisably different. In some
cases of change it is easy to observe the difference in pro-
perties. For example, if we put a piece of paper into a flame,
the paper burns and chars : from being white coherent paper it
becomes a brittle black mass of obviously different properties.
On the other hand, there are many cases of change where the
difference in properties of the changed and unchanged sub-
stances is by no means easily detected by simple inspection.
A piece of quicklime is not unlike the piece of limestone from
which it was derived by heating, but if we observe its properties
more closely, we find that it must be regarded as an entirely
different kind of substance. One important respect in which
it differs from limestone is that when treated with a small
quantity of water, it becomes warm, cracks, and breaks up into
a bulky friable mass. The original limestone exhibits none of
this behaviour, being practically unaffected by water. We can
use, then, the action of water as a test for distinguishing be-
tween limestone and quicklime. Chemical tests are in general
of this nature : if we cannot tell offhand what a substance
is, we try its action with something else, and see how it
behaves.

[ INORGANIC CHEMISTRY

Quicklime is produced from limestone on the lai^e scale
n the operation known as lime-lmrnuig. The operation is
- conducted in a lime-kiln, which
is a kind of furnace of stone or
brickwork, the internal cavity
being usually either conical or
egg-shaped. This cavity is
filled with alternate layers of
fuel (coal or turf) and lime-
stone, air being admitted at
the bottom. When the fuel
burns in the air, the limestone
is raised to such a temperature
that it is converted into quick-
lime. The quicklime falls as
the fuel burns away, and may
be raked out through the air-
holes, while fresh fuel and
limestone are supplied from
above, to take its place, the
- process thus being made to

Fig. I. — ^Kiln for Lime- Burning.

go on continuously.

When limestone is burnt in
Ri«^i'f?ay«s of' fud ^d'li'^siol,* this way it loses nearly half its

joverC. The gK« «cape by the f
The con Lea! grating C directs the pieces
burnt linie >o a> lo faciliuti withdraw

p by lemoyini! tie weight, ooly about 1 1 cwts.
of quickhme bemg obtamed
from a ton of limestone. It is
clear that, to justify this loss of
material, and the loss of fuel in producing it, the quicklime
must possess some valuable property not possessed by the
original limestone. The property in question is the power
of being slaked by water. As has already been indicated,
quicklime, when brought into contact with a small amount
of water (about one-third of its weight), swells and crumbles,
the mass becoming hot at the same time. The heat produc-
tion may be so great that some of the water is converted
into steam. The powder obtained from quicklime and water
is called slaked Ume, and its properties are plainly different
from those of the substances from which it was produced. It
bears no resemblance whatever to water, and differs from
the quickhme used in its formation in being unaffected by

EXAMPLES OF CHEMICAL ACTION 3

further amounts of water. No heat is produced when these
substances are brought together, and the slaked lime is merely
wetted by the contact, or dissolved away if a very large
quantity of water is used. A paste of slaked lime and
water is used in the preparation of mortar for building pur-
poses, and it is to this end that limestone is chiefly burnt.
The slaking of quicklime for mortar or plaster may be seen
during the construction of almost any house, the materials
being mixed in large wooden troughs, from which clouds of
steam rise as the water is added.

Definite Weights. — If we take a portion of pure limestone,
say 100 grams, and subject it in a crucible to a bright red
heat, it will speedily begin to lose weight, but the loss of
weight will not go on indefinitely. After the weight of the
substance has been reduced to 56 grams no further loss occurs,
however long the heating may be continued. The limestone
has then been entirely converted into quicklime, which is quite
unalterable by heat. Now, we shall always find that if we
convert a pure limestone into quicklime, 100 parts of the
limestone will give 56 parts of quicklime \i,€, about 11 cwts.
per ton), no matter how the burning may have been con-
ducted. Here, then, we have a definite and fixed proportion
between the weight of limestone and the weight of quick-
lime derived from it; and we shall find that such definite
relations by weight are characteristic of all chemical chsCnges.

Should we, however, stop heating before the limestone has
lost weight to the full extent — suppose, for example, that the
100 parts have become 78 parts — of what substance does the
residue then consist ? We have stopped the process when the
loss of weight is only one-half the total possible loss — i,e, 22
parts instead of 44. Is the substance left behind exactly
intermediate in properties between limestone and quicklime ?
In a certain sense it is, but on closer inquiry we find that the
residue is not a single substance at all, but is really a mixture
of two substances — namely, unchanged limestone and quick-
lime. This may be shown in various ways. For example,
if we treat the mixture with water, the limestone half will be
entirely unaffected, while the quicklime half will be slaked ;
and the slaked lime so produced may be entirely dissolved
away if a sufficient quantity of water is taken.

4 INORGANIC CHEMISTRY

There is thus really no intennediate step in the change
from limestone to quicklime. Any limestone which has
undergone alteration has been completely changed into quick-
lime: the rest is limestone with all its original properties.
This sharp passage of one substance into another, as dis-
tinguished from a gradual transition, is always found in
chemical change.

If we inquire as to the reason why limestone loses weight
when converted into quicklime, we find it to be a well-known
fact that the atmosphere in the immediate neighbourhood of
limekilns is dangerous, inducing sleep and even causing death
by suffocation. This is due to the presence near the kilns of a
heav)* vapour given off by the limestone during the burning, and
usually called carbonic acid gas. If all the carbonic acid
given off by the limestone on its conversion into quicklime
were collected and weighed, its weight would be found to be
exactly equal to the loss in weight experienced by the lime-
stone — 1>. 44 parts for each loo parts of the original lime-
stone.

Consider now the slaking of lime. Again we meet with
definite weights. The 56 grams of quicklime derived from
100 grams of limestone will only take up 18 grams of water
to form dry powdery slaked lime. If we add less than 18
grams of water, the substance obtained is a mixture of
quicklime and slaked lime ; if we add more than 18 grams, a
mixture of slaked lime and water remains — ue, a wet mass
or paste instead of the dry powder. There is thus once
more a sharp transition of one substance into another of utterly
different properties, no intermediate stages being formed.

We may now write the chemical actions involved in the
burning and slaking of lime as follows : —

Limestone j^ives Quicklime and Carbonic acid gas
100 parts = 56 parts + 44 parts

Quicklime and Water give Slaked Lime
56 parts + 18 parts = 74 parts

These instances exemplify the fact that there is no loss of
weight in chemical action if all the substances involved in the
action are taken account of. The apparent loss of weight
in the burning of limestone is due to the carbonic acid gas

EXAMPLES OF CHEMICAL ACTION 5

being permitted to escape without being weighed. When
it too is weighed, the sum of its weight and that of the
quicklime is exactly equal to the weight of the limestone
taken. Again, the weight of the slaked lime obtained from
quicklime and water is exactly equal to the sum of the weights
of these substances. Though, therefore, there may be great
alterations in the properties of substances after a chemical
change, there is no alteration in the total weight. This has
been found to be strictly true for every chemical action that
has been accurately investigated.

Varieties of the same substance. — There are other sub-
stances besides limestone which on heating become converted
into quicklime and carbonic acid gas. Such are chalk, white
marble, precipitated chalk, and Iceland spar. These substances
are easily distinguished from limestone and from each other
by even a superficial examination. Chialk is, as a rule, much
softer than limestone, marble much harder and capable of
acquiring a fine polish; Iceland spar forms perfectly clear
transparent crystalline masses; precipitated chalk is a soft
white powder. Yet if pure specimens of these bodies are
heated, we find that in each case 100 parts of the substance
will give 56 parts of quicklime and 44 parts of carbonic
acid gas; and the quicklime produced by their calcination
will unite with 18 parts of water to give slaked lime of exactly
the same properties, whether it is prepared from chalk, marble,
Iceland spar, or limestone.

We have here, then, five apparently different substances
which behave chemically towards heat in precisely the same
manner. These substances are said to be chemically identical,
and are spoken of as different varieties of the same chemical
substance. A greyhound and a bull-dog are very different in
superficial appearance, yet they are both classed as varieties
of dog, because they answer to certain zoological tests which
are used to identify the species dog. In the same way lime-
stone, marble, chalk, etc., are all classed as varieties of the
chemical substance calcium carbonate, because they answer
to certain chemical tests by means of which calcium carbonate
may be identified — in particular, because they split up into
quicklime and carbonic acid gas in the proportions by weight
that have been indicated above.

6 INORGANIC CHEMISTRY

diemical Nomenclatiire and Equations. — Just as zoologists
have found it convenient for systematic purposes to give
animals names which differ from the common names of those
animals — using for example Cams familiaris instead of dog —
so chemists have devised a systematic nomenclature for sub-
stances regarded in the purely chemical aspect, with the
result that one and the same thing has very frequently two
names, one the common name, the other the systematic
chemical name. Thus we have —

Common Name Systematic Chemical Name

Marble, Chalk, Limestone, etc. Calcium carbonate

Quicklime Calcium oxide

Carbonic acid gas Carbon dioxide

Slaked lime Calcium hydroxide

The burning of limestone, chalk, or marble, is, then, chemically
speaking, one action which may be represented as follows : —

Calcium carbonate gives Calcium oxide and Carbon dioxide
I GO parts = 56 parts + 44 parts

Similarly for the slaking of lime we may write —

Calcium oxide and Water give Calcium hydroxide
56 parts + 18 parts = 74 parts

Such brief statements of chemical actions as those just given

are called chemical equations, the total weights on the tw^o '

sides being equal. As we shall see later, it is not customary

to write chemical equations at length, giving both names and

weights as above : chemical formulae are used instead of the

chemical names, and from the formulae the weights can be

calculated.

It should be noted that equations have only validity for
pure chemical substances. Thus a limestone on complete j

burning may be found to give more or less than 56 parts of j

burnt lime for 100 of the original limestone. Such limestone,
however, is impure — that is, the calcium carbonate of which
the bulk of it consists is mixed with smaller quantities of
other substances. The chemical equation in this case only »

applies to the calcium carbonate part of the impure limestone,
nothing being stated as to how the impurities are affected by
the burning.

EXAMPLES OF CHEMICAL ACTION 7

The two chemical actions which we have just considered
belong to two different and important types. The process
of lime-burning is a chemical decomposition, or the splitting
up of a substance into two or more others. The process of
lime-slaking is a chemical combination, or the union of two
or more substances to form a single chemical compound. We
shall meet in the sequel with many examples of both of these
types of chemical action.

CHAPTER II

CONDITIONS OF CHEMICAL ACTION

In the last chapter we saw under what conditions calcium
carbonate is decomposed into calcium oxide and carbon
dioxide : in order to effect the decomposition, the substance
must be raised to a red heat. At the ordinary temperature
calcium carbonate, if left to itself, undergoes no change; or, as
chemists are accustomed to say, it is under these conditions
perfectly stable. The high temperature is necessary before
any decomposition occurs. This behaviour is very general ;
most substances which are stable at the ordinary temperature
decompose when heated to a sufficiently high temperature.
Sugar, for example, chars when heated, and gives off vapours,
the decomposition being accompanied by the familiar smell
of burnt sugar. Wood, too, if heated by itself, is converted
into combustible vapours and charcoal. An instance of
decomposition by heat very similar to that of limestone is
afforded by slaked lime. If slaked lime be raised to a red
heat, it decomposes into quicklime and water according to
the equation —

Calcium hydroxide giv Calcium oxide and Water.

74 parts = 56 parts + 18 parts

the water at that temperature being in the form of highly
superheated steam. A high temperature then is generally
favourable to decomposition.

If we now consider what conditions are favourable to
chemical combination, or in general, chemical action between
two or more substances, we see in the first instance that
the substances must be in contact with each other before they
can interact at all. So long as they are separate, chemical
action between them is impossible. Quicklime and water
must be brought together before slaked lime can be produced
by their union, It would appear reasonable, therefore, that

CONDITIONS OF CHEMICAL ACTION 9

if we wish to promote chemical action between two substances,
we should bring them into very <:lose contact — i.e. mix them
as thoroughly as possible: Large pieces of solid substances
when brought together have only a very small surface of
contact, but if each solid is reduced to a powder, the exposed
surface is very great, and when the two powders are mixed,
the surface of contact between them is much increased. Thus
the charcoal, sulphur, and nitre which are the ingredients of
gunpowder, would not form an explosive if mixed together in
pieces the size of a pea. The surface of contact between the
pieces would not be sufficiently great for that rapid chemical
action to take place on which the explosion of gunpowder
depends. In the manufacture of gunpowder, each ingredient
is ground separately to a fine powder, and the powders are
then intimately mixed and caked together by means of a
small quantity of water, the cake being afterwards dried and
broken up into grains of the requisite size. Each grain then
contains all the ingredients necessary for the rapid chemical
action brought about when the gunpowder is fired.

If vigorous action is desired between a solid and a liquid,
the solid should be powdered before being brought into
contact with the liquid. This may be readily seen by noting
the vigour of the action between quicklime and water when
the quicklime is finely powdered. Even when the quicklime
is in lumps, however, the action is still vigorous, because the
quicklime is somewhat porous and soaks up the water into the
interior of the lumps.

Solntion. — A method which chemists very frequently employ
for bringing about the very intimate contact of substances which
are ordinarily solid or liquid, is to dissolve them in water and
then mix the solutions together. In each solution the dissolved
substances are distributed quite uniformly through the liquid,
so that when the two solutions are mixed together, which can
be done very simply and thoroughly by stirring, the dissolved
substances are brought into a contact much closer than could
generally be obtained by reducing them to a powder and then
mixing. When water cannot be used as a solvent, some other
liquid, such as spirit or ether, may be employed ; but in the
laboratory, as in common life, water is used for the purpose in
ninety-nine cases out of a hundred.

10 INORGANIC CHEMISTRY

Temperature of ReactioiL — Some substances act at once
when brought into contact. Thus when quicklime and
water are placed together, they immediately unite to give
slaked lime. On the other hand, some substances which
are capable of entering into vigorous chemical action do
not affect each other under ordinary conditions, however
long they may remain in contact. Thus the nitre, sulphur,
and charcoal of gunpowder are absolutely without action on
each other until the temperature is raised to a sufficient
degree, when the ingredients suddenly react with such
violence as to produce an explosion. Coal-gas and air,
again, may be mixed at the ordinary temperature without
anything occurring, but if a light is applied chemical action
may take place with explosive violence. That a high tem-
perature is necessary for the interaction of air and coal-gas
may be easily verified by attempting to light a jet of the
gas by means of a red-hot poker. Notwithstanding the high
temperature to which the gas in contact with the poker is raised,
it will not ignite; but as soon as a lighted taper is applied,
ignition follows, for the temperature of the lighted taper is
very much higher than the red heat of the poker.

If such a high temperature is necessary for the ignition of
coal-gas, the student may ask: "Why does not the gas go
out when the flame which lit it is removed ? " The reason is
simple. Vigorous chemical action is almost invariably accom-
panied by production of heat, and the heat produced by the
chemical action of the coal-gas and air is sufficient to keep
the temperature far above the ignition-point once the action is
started. The case of gunpowder is similar. It is not necessary
to heat the whole mass of the powder in order to fire it ; if a
light is applied to one portion, that is sufficient. The heat
given out by the first portion as it explodes serves to raise the
temperature of neighbouring portions to the exploding point,
and so the explosion is propagated through the mass.

Some actions do not produce sufficient heat to keep them-
selves going once they are started. Thus a jet of ammonia
gas will burn at the flame of a taper so long as the taper is
kept at the jet, the heat from the lighted taper being sufficient
to raise the mixture of ammonia gas and air to the ignition
point. But as soon as the taper is removed, the action ceases,
for the burning of thQ g^mnaonia dioes not; of itself produce

CONDITIONS OF CHEMICAL ACTION ii

enough heat to keep the temperature of the ammonia and
air up to the point at which they interact.

It may be said in general that chemical actions take place
more rapidly as the temperature is raised. This circumstance
explains the increase in vigour observed in many chemical
actions which are at first comparatively slow. When cold
water is poured on a lump of quicklime, nothing apparently
happens for the first moment; but very soon a little steam
makes its appearance, the action gets brisker, the evolution of
steam becomes more rapid, and the mass cracks and crumbles
to the powder of calcium hydroxide. This increase in vigour
is chiefly due to the rise in temperature caused by the inter-
action of the first portions. The temperature of the neigh-
bouring portions of water and quicklime is raised, with the
result that they react much more rapidly than the first portions,
and so produce more heat, which increases the rate of reaction
of remaining portions, a constant acceleration of the reaction
thus going on.

If the temperature at which the quicklime and water are
brought into contact is above loo degrees, the water is in the
form of water-vapour or steam, but the action goes on just as
before. Should the quicklime and water-vapour, however, be
raised to a bright red heat, the action ceases, and no union
takes place at all. This is in contradiction to the statement
made above, that rise of temperature increases the vigour of
chemical action, but a closer consideration shows that this
exception is more apparent than real. We have already seen
that calcium hydroxide decomposes at a red heat into calcium
oxide and water-vapour. Now this is exactly the reverse of
the reaction we have been considering, so that if calcium
hydroxide were formed at the high temperature, it would at
once be decomposed into the substances from which it was
produced. We are, in fact, here dealing with a reversible
action, the equation of which is —

Calcium oxide + Water ^ Calcium hydroxide
56 + 18 = 74

The oppositely directed arrows used instead of the sign of
equality indicate that the action can proceed backwards as
well as forwards. The forward action is a combination, the

12 INORGANIC CHEMISTRY

reverse action a decomposition. Rise of temperature may
increase the vigour of both reactions, but it favours the
decomposition at the expense of the combination, so that at
a bright red heat the direct action is altogether overpowered
by the reverse action.

A similar instance of opposed reactions is afforded by
the burning of limestone. At a bright red heat the action is
one of decomposition, the calcium carbonate splitting up into
calcium oxide and carbon dioxide. Carbon dioxide, at a
lower temperature, is re-absorbed by calcium oxide with for-
mation of calcium carbonate, so that we have the reversible
action —

Calcium carbonate T^ Calcium oxide + Carbon dioxide.
lOO =56 + 44

Rise of temperature favours the decomposition, fall of tem-
perature the recombination.

It must not be supposed that all chemical actions are
reversible; indeed, most of the actions that the elementary
student of chemistry encounters are irreversible. The charcoal
and combustible vapours obtained by heating sugar will not
recombine on cooling to reproduce the sugar from which they
were formed, nor will the ash and gases from gunpowder
which has been fired ever interact to form the original mixture
of charcoal, sulphur, and saltpetre.

CHAPTER III

SOLUTIONS AND SOLUBILITY

It has just been stated that chemists are in the habit of
dissolving solid substances in liquids, especially water, in
order to obtain them in a state suited to the production of
many chemical actions. It is usually said in this connection
that when a substance is dissolved in water, its chemical
properties are not greatly changed; but this statement must
not be taken too literally, and due allowance must be made
for the properties of the water in the solution. For instance,
a lump of sugar will bum if a light is applied to it; but if
the same lump is dissolved in water we can scarcely expect
it to bum then, were it only for the plain reason that the
burning sugar would be extinguished by the water. We
shall see in the sequel that water profoundly modifies the
chemical properties of many other substances, such as nitric
and sulphuric acids, and it will be necessary for us to know
the properties of aqueous solutions of these substances, as
well as those of the pure substances themselves, since the
former are often of equal if not greater importance than
the latter.

The solvent action of water varies very much with the
substance on which it acts. Thus pure water will scarcely
dissolve any calcium carbonate, whether in the form of chalk,
marble, or limestone, whilst it will readily dissolve large
quantities of sugar, salt, or nitre. To ascertain the extent
of the action of water on a solid substance, the solid should
be finely powdered and shaken up for a long time with the
water, in order to get thorough contact between them. After
a time (usually several hours) the water will take up no more
of the solid, and is then said to be saturated with it. The
saturated solution is then decanted or filtered off from the
excess of solid, and the water driven off from the solution
at a gentle heat. If the quantity of solution which is evapor-

INORGANIC CHEMISTRY

ated, and the residue which is derived from it, are both
weighed, we can state in numbers the extent of the solvent
action of the water, or, what is the same thing looked at
from another point of view, the solnbility of the salt in water.
Thus we find that at the ordinary temperature the quantities
of water necessary to dissolve one part of substance are as
follows : —

Substance

Parts water

Cane sugar ....

Salt . .

Saltpetre (nitre)

Slaked lime ....

Calcium carbonate

4
800

I 000000

To express the same results in another way, we may say
that the following quantities of the various substances are
dissolved by 100 parts of water : —

Substance

Solubility in 100 parts
of water

Calcium carbonate .

O'OOI

Slaked lime .

013

Saltpetre (nitre)

Salt

. . . 36

Cane sugar .

200

It is usually in this second way that solubilities are now stated.

From these tables it appears that cane sugar is extremely
soluble in water, salt and saltpetre freely or easily soluble,
calcium hydroxide sparingly soluble, and calcium carbonate
practically insoluble. The solubility in water of a sub-
stance such as sulphur is too small to be measured, and
it is therefore said to be insoluble in water.

We generally find that warm water dissolves things better
than cold water, the difference in solubility being sometimes
very great. Thus 100 parts of water at the boiling point
(loo**) will dissolve 39 parts of salt, and 250 parts of
nitre, instead of 36 and 25 parts respectively at the ordinary
temperature. The solubility of salt is therefore only slightly
increased by the rise of temperature, but the solubility of the
saltpetre is ten times as great in boiling water as in cold
water.

SOLUTIONS AND SOLUBILITY

The variation of the solubility of substances with change
of temperature is conveniently represented by means of a
solubility diagram (fig. 2). Temperatures are measured
along horizontal lines, and the amounts of substance dis-
solved along vertical lines. The solubility of the substance

240

10*' 20° 30° 40" '60" 60** 70** 80° 90" lOO'

Temperature

Fig. 2. — Diagram showing Variations of Solubility with

Temperature.

at any one temperature is thus represented by a point in
the diagram, the position of which indicates on the horizontal
scale the particular temperature considered, and on the
vertical scale the number of parts of substance dissolved by

1 6 INORGANIC CHEMISTRY

loo parts of water. When these points for all the different
temperatures are joined up, a solubility curve is obtained.
Each of the substances considered in the diagram has its
own solubility curve. The more nearly horizontal the curve
runs {e,g, salt), the smaller is the variation with the tempera-
ture ; the steeper the curve is {e.g, nitre), the greater is the
variation with temperature. When the curves for two different
substances cut, it shows the substances are equally soluble
at the temperature represented by the point of intersection.
Thus salt and nitre are shown by the diagram to have the
same solubility (namely, 36 parts in 100 of water) at 24**.

Occasionally it happens that the solubility of a substance
is less in hot than in cold water. The solubility of calcium
hydroxide in boiling water is 0.06 — i.e. only half the solubility
in cold water. This diminution of solubility with rise of
temperature can be easily shown experimentally as follows : — If
the clear saturated solution of calcium hydroxide, commonly
known as "lime water," is heated in a glass vessel to the
boiling point, it is seen to become turbid, owing to separa-
tion of solid calcium hydroxide. The water at 100**
being able to hold only half as much calcium hydroxide
in solution as it could at the ordinary temperature, deposits
the other half in the form of white solid particles which render
the solution milky.

Since nitre is ten times as soluble at 100° as it is at the
ordinary temperature of 15°, a solution saturated at the boiling
point and then cooled will part with nine-tenths of the nitre
it held dissolved. The nitre falls out from the solution if the
cooling is rapid in the form of gritty particles technically
known as " nitre meal." These particles are in reality small
crystals : and if the hot solution is allowed to cool very slowly,
large crystals of nitre may be obtained.

Nearly all chemical substances are crystalline, or may be
made to assume the crystalline state. Crystalline substances,
when properly investigated, are found to possess a definite
form or shape of their own : non-crystalline or amorphous
substances are formless. Sugar is an example of a substance
which crystallises well, and nearly all the sugar used nowadays
is in the form of small crystals. Clusters of large crystals of
sugar grown together may be seen in sugar-candy. If a piece
of sugar-candy is examined, it will at once be evident that the

SOLUTIONS AND SOLUBILITY

Crystal of Cane>Sugar.

same shape is repeated over and over again, and that the
surfaces of the separate crystals are plane. This is character-
istic : the faces of all perfect crystals are
perfect planes. Frequently crystals are so
small that the eye cannot tell if they possess
any regular shape or if they are bounded
by plane surfaces. In such a case recourse
may be had to the reflection of light.
Plane surfaces reflect light better than
irregular surfaces, so that the small crystal
faces may be made to appear as bright
specks when they are held in the proper
position with respect to the eye and a
source of light. The small crystals of
which a piece of lump sugar consists may
easily be detected in this way if the lump
is held so that the light from a window
falls upon it and is reflected upwards to
the eye.

Solids may almost always be made to
crystallise either by fusing them and allowing them to solidify,
or by dissolving them in some solvent and letting them
separate from solution. If the separation is rapid, as in the
above instance of nitre meal, the crystals are invariably small.
^Vhe^ large crystals are desired, the separation of the substance
must be allowed to go on slowly without much mechanical dis-
turbance of the solution. This may be effected either by very
slow cooling, or by letting the solvent evaporate slowly from
the saturated solution. Good crystals may be obtained by
allowing a cold solution of photographic hypo (sodium thiosul-
phate) to stand in an open vessel, the solvent being gradually
lost by evaporation.

The process of crystallisation from solution in this way is a
very important one, as it enables us to remove the impurities
from impure substances. For example, if we grind together
some alum and a little . blue vitriol (copper sulphate) in a
mortar, the mechanical separation of the two substances in
the mixture is practically impossible. We can remove the
copper sulphate, however, and obtain crystals of pure alum,
by dissolving the mixture in hot water and allowing the solu-
tion to crystallise. The alum crystallises^ out in perfectly

i8 INORGANIC CHEMISTRY

colourless crystals, which are quite free from copper sulphate.
Even if copper sulphate crystallises out at the same time as
the alum, it will form separate crystals, which if the crystallisa-
tion takes place with the necessary slowness, may reach such
a size as to be easily removed from the alum crystals by
hand.

Some substances, when they separate from aqueous solution,
do so as crystals which contain a definite amount of water.
Thus, if we dissolve pure copper sulphate in water and allow
the solution to crystallise, the substance which separates is not
entirely copper sulphate, but a substance (blue vitriol) which
contains copper sulphate and water in the proportions of i6i
parts of the former to 90 parts of the latter. The water in
such a case has none of the properties of liquid water, and is
called water of crystallisation.

CHAPTER IV

STMBOLS—FOBMULJE— EQUATIONS

In the splitting up of calcium carbonate by heat we have seen
an instance of chemical decomposition. Similar decomposi-
tion is possible for most chemical substances, by heat or some
other agency, and the question arises : Is there any limit to
chemical decomposition ? Are all substances decomposable ?
As an answer to these questions chemists have found by
experiment that there are some seventy substances which have
resisted all attempts to decompose them. These are called
simple substances or elements. All other substances, which
are known as compounds, are made up of these elements, and
can be decomposed either directly or indirectly.

We can most readily express the composition of any
chemical compound by stating of what elements it consists,
and the quantities of these elements that are combined
together to form the compound. Now, it has been found that
without exception pure chemical substances have an invariable
composition — ue. they are made up of the same elements
combined together in the same proportions, no matter what
the source of the compound may have been. For example,
calcium carbonate, whether it is artificially prepared, or has
been found in nat<ure as marble, Iceland spar, etc., always
contains the elements calcium, carbon, and oxygen; and its
invariable composition is

Calcium 40 parts

Carbon . . . . 12 „
Oxygen 4^ »

100 parts

We can express this and similar compositions very simply
by making use of a system of symbols which chemists have
invented for the purpose. A list of the commoner elements

INORGANIC CHEMISTRY

considered in this book is given below. The list will be
observed to contain all the common metals.

Table of Commoner Elements

Name

Symbol

Weight

Aluminium .

Barium

137

Bromine

Calcium

Carbon

Chlorine

35-5

Copper

Gold (Aurum)

197

Hydrogen

Iodine

127

Iron (Ferrum)

Lead (Plumbum) .

207

Magnesium . . . .

Mercury (Hydrargyrum)

200

Nitrogen . . . .

Oxygen . . . .

Phosphorus

Potassium (Kalium)

Silver (Argentuui)

108

Sodium (Natrium)

Sulphur . . . .

Tin (Stannum)

118

Zinc . . . . .

65-5

For each element there is a symbol, which consists of the
first letter of the Latin name of the element, together some-
times with one of the subsequent letters, in order to prevent
confusion when the names of several elements begin with
the same letter. By writing these symbols alongside each
other, we can easily express what elements any given com-
pound contains. Thus, to express that calcium carbonate
contains calcium, carbon, and oxygen, we have only to write
CaCO. But besides being mere shorthand for the names of
the elements, these symbols are something more. Each
symbol expresses a definite amount of the element which

SYMBOLS— FORMULAE—EQUATIONS 2 1

it denotes. These combining weights are given in the table
after the symbols. Thus the symbol C not only indicates
the element carbon, but 1 2 parts by weight of carbon : the
symbol Ca represents not only calcium, but 40 parts of calcium,
and so on. The complex symbol or formula CaCO represents,
then, not only a compound containing the elements calcium,
carbon, and oxygen, but a compound whose composition is

Calcium ..... 40 parts

Carbon . . . . . 12 „

Oxygen . . • . . . 16 „

68 parts

This compound, which contains 40 parts of calcium in 68,
cannot be calcium carbonate, which contains 40 parts of
calcium in 100. If we compare the compositions

CaCO

Calcium carbonate

Calcium .

Carbon .

Oxygen .

we see that relatively to the other elements, calcium carbonate
contains three times as much oxygen as a compound of the
formula CaCO. We may therefore write its symbol CaCOOO,
which now expresses the correct composition.

In order to save repetition of symbols, it is customary to
write the formula of calcium carbonate CaCO 3 instead of at
length as above. The number affixed to the symbol of an
element indicates how often the symbol must be repeated.
Thus the formula of sulphuric acid H2SO4 is a shorter form
of HHSOOOO. This formula expresses the fact that sulphuric
acid contains

Hydrogen . H2= 2x1= 2 parts
Sulphur . S = 32 = 32 „
Oxygen . O4 = 4x16 = 64 „

98 parts

In the case of calcium carbonate, the formula of the substance
expresses the percentage composition directly, as may be seen

22 INORGANIC CHEMISTRY

above, but this is merely a coincidence. In general the formula
does not give the number of pavts of each element in loo
parts, but in some other number. For sulphuric acid the
formula expresses the number of parts of each element in 98.
It is, of course, easy to calculate from this by simple proportion
the number of parts in 100, or the percentage composition :

Parts in 98 Parts in 100

Hydrogen. . . 2 2X-— = 2.04

100
Sulphur ... 32 32 X— ^=32.65

Oxygen ... 64 64 x— — = 65.31

98
100

IOC

Similarly, if we are told that the formula of sodium chloride is
NaCl, we can calculate its percentage composition thus :

100
Sodium . . Na=23 23 x -g-p = 39.32

Chlorine . . CI = 35.5 35.5 x -—— = 60.68

Per cent.
100

58.5

In using these formulae, the student must clearly understand
that the composition of any chemical compound must be
ascertained by actual experiment before a formula for it can
be written at all. The formula is merely a brief and convenient
method of expressing the experimental results, and must never
be conceived by the beginner in any other sense.

It has already been indicated that chemical formulae are
used in equations instead of the names and weight of the
substances involved. Thus the equation

Calcium carbonate = Calcium oxide + Carbon dioxide
100 56 44

can be expressed in formulae as follows: —

CaCOa = CaO + CO 2

SYMBOLS— FORMULA—EQUATIONS 23

This symbolic equation gives us at once the correct weights, if
we refer to our table of symbols :

CaCOg = CaO + CO2

40+12+48 40+16 12 + 32

100 = 56 + 44

The equation, of course, presupposes that we know the com-
position of calcium oxide and of carbon dioxide to be
represented by the figures corresponding to the above formulae.
Similarly, if we know the composition of water and of
calcium hydroxide to be given by the formulae H3O and
CaH202, we can write the equation for the slaking of lime
as follows : —

CaO + H2O = CaHgOo

Reference to the weights expressed by the symbols will show
that this equation indicates the combination of 56 parts of
calcium oxide with 18 parts of water — ue. the proportions
actually found by experiment.

We can always test if an equation is arithmetically correct
by adding up the weight values on the right and left of the
equation, and seeing if they are in reality equal. A simpler
plan of doing this, however, is to count the number of symbols
of each element on the two sides. If the equation is
arithmetically correct, the symbol of each element must
appear the same number of times on the two sides, for it
represents a fixed quantity of the element, and the elements
are not transformable into one another. In the equation for
the slaking of lime, we have on each side, one Ca, two O's,
and two H's : the equation is therefore arithmetically correct.
The testing of the arithmetical accuracy of an equation is very
important, for the reason that an arithmetically incorrect
equation cannot by any possibility be chemically correct. For
an equation to be chemically accurate, the symbol of each
element must appear the same number of times on the two
sides, otherwise the total weight of the substances involved in
the action would be changed, or else one element would have
undergone transformation into another, both of which assump-
tions are chemically speaking impossible.

Of course, and the student must have this constantly before
him, it does not follow that an equation which is arithmetically

24 INORGANIC CHEMISTRY

correct is also chemically correct — ue, that it is an expression
for a chemical action which actually occurs. For example, the
equation,

CaCOa = CaOa + CO

Calcium carbonate Calcium peroxide Carbon monoxide

is arithmetically accurate, and, moreover, involves only sub-
stances which are actually known; yet it is chemically quite
inaccurate, for by no means at present known can we decom-
pose calcium carbonate into calcium peroxide and carbon
monoxide.

Before we can write an equation which shall be chemically
accurate, we must know then not only the composition of all the
substances involved, but also that the change implied by the
equation actually takes place. A chemical equation ought
always to be the expression of a chemical fact, and what is
fact can only be ascertained by trial — i,e, by experiment.

The student will therefore do well to remember that though he
can always writ» equations which are arithmetically accurate —
an arithmetically inaccurate equation can only be the result of
carelessness, for the accuracy can always be tested by counting
the symbols — he cannot, without a real knowledge of \}[i^ facts
of chemistry, write equations which are chemically accurate.

There are one or two points concerning the use of figures in
connection with chemical symbols with which the student must
make himself familiar. The formula CO 2 expresses 44 parts of
carbon dioxide. Now, suppose we wish to write a formula ex-
pressing three times as much as this — viz, 132 parts. This can
be done in two ways — either we can write CsOg, which still
preserves the proper proportion between carbon and hydrogen,
and trebles the amount, or else we write 3CO2, the number
prefixed to the formula applying to all the symbols in it. The
latter method is almost invariably adopted by chemists, who
prefer to keep the actual formula as simple as possible. The
formula HgO stands for 18 parts of water; 2H2O stands for
36 parts; 3H2O for 54 parts, and so on. It should be
noted that fractional numbers never appear in chemical for-
mulae or equations.

Sometimes we find it convenient to group certain symbols
together within a formula. Thus calcium hydroxide, which
we have written CaHgOo, is usually written Ca(0H)2 for a

SYMBOLS— FORMULAE— EQUATIONS 25

reason which will appear in the sequel. A number affixed
in this way to elements within brackets applies to all the
symbols contained in the brackets : Ca(N03)2 is the same as
CaNgOe ; (NH4)2S04 is the same as N2H8SO4. Occa-
sionally we find symbols in a complex formula separated
by a point or a comma. The formula CuS04,5H20 or
CUSO4.5H2O is the formula of blue vitriol (p. 18), and indi-
cates that the substance contains 159 parts of pure copper
sulphate, together with 5 x 18 = 90 parts of water of crystallisa-
tion. If a number is prefixed to such a complex formula it
only applies up to the point or comma. Thus 2CuS04,5H20
is not twice the formula of blue vitriol, but the formula of a
substance containing 2 x 159 parts ofcopper sulphate and 5 x 18
parts of water. If we wish to double the whole of the complex
formula we put it all within brackets, and then prefix 2, thus :
2(CuS04,5H20), which indicates two formula-weights of blue
vitriol.

As an example of equation writing we may take the follow-
ing : — It is known that when sodium hydroxide and sulphuric
acid react under certain conditions, sodium sulphate and water
are the only products. It is further known that the composition
of these substances may be accurately expressed by the follow-
ing formulae : —

Sulphuric acid .... H2SO4

Sodium hydroxide .... NaOH

Sodium sulphate .... Na2S04

Water HgO

Knowing these facts, we can proceed to write the equation as
follows. First we write down the formulae joined by the usual
algebraic symbols,

H2SO4 + NaOH = Na2S04 + HgO

This is plainly not an equation, as we can see by counting the
symbols on the two sides. On the left we have Na represent-
ing 23 of sodium ; on the right we have Na2 representing 46
of sodium. Since only whole numbers must appear in chemical
equations, we cannot halve the amount of sodium on the right,
for that would entail halving the single symbol S in order to
keep the composition of sodium sulphate. We must therefore
double the amount of sodium on the left, which we can only

26 INORGANIC CHEMISTRY

do by doubling the whole formula of sodium hydroxide. We
thus get —

H2SO4 + 2NaOH = Na2S04 + H^O

This is still not an equation, for we have four H's and two O's
on the one side, as against two H*s and one O on the other.
To remedy this we double the quantity of water on the right,
and so obtain the real equation —

H2SO4 + 2NaOH = Na^S04 + 2H2O

in which the symbols are properly balanced. This equation is
arithmetically correct, and tells us that 98 parts of sulphuric
acid react with 80 parts of sodium hydroxide to give 142 parts
of sodium sulphate and 36 parts of water. Here, by knowing
the formulae of the reacting substances and the products of the
reaction, we have been able to calculate the proportions in
which all these substances are involved in the chemical action
by merely getting the chemical equation to balance. This is
in general the case ; if we know all the reacting substances and
all the products of the reaction, and, further, know the formulae
which express their composition, we can, by a purely arith-
metical process, write an equation which expresses the propor-
tions by weight in which all the substances concerned partici-
pate in the reaction. It is sometimes not easy to arrive at the
arithmetical solution of an equation in the manner indicated
above, but it is always possible, and only requires a little
expertness, which comes of practice. Examples of various
methods of solving will be given as occasion requires.

In ordinary chemical equations the reacting substances are
put on the left, and the products of reaction on the right. It
is therefore necessary always to read chemical equations from
left to right, unless the action is shown to be reversible by
means of oppositely directed arrows. Thus the reversible
equation

CaCOs '^ CaO +- CO2

really gives expression to the two ordinary equations —

CaCOa = CaO 4- CO2
CaO + CO3 = CaCOs

each of which is the reverse of the other.

CHAPTER V

COMBUSTION

There are a great many chemical reactions in which air plays
a part. Atmospheric air consists of about one measure of
oxygen and four measures of nitrogen. Although oxygen
therefore occupies only one-fifth of the total volume of the
air, it is nevertheless the active component of the mixture,
the nitrogen merely serving to moderate the activity of the
oxygen.

Familiar instances of actions in which air takes part are to
be found in the processes of combustion. For example, when
fuel such as wood or coal burns in the air, the chief action
is the union of the components of these substances with the
oxygen in the air, the combustion taking place with such
vigour that heat and light are developed. Similarly, the
burning of a taper or coal-gas in air consists in the union of
the materials of these substances with oxygen. The principal
components of such ordinary fuels and illuminating agents are
the elements carbon and hydrogen. On combustion the carbon
unites with the oxygen to form carbon dioxide, and the
hydrogen unites with oxygen to form water, according to the
equations —

c +

CO,

Carbon

Oxygen

Carbon dioxide

2Hj +

iK^O

Hydrogen

Oxygen

Water

That water is produced in the form of vapour may be easily
shown by holding a cold bright object such as a polished
piece of metal a little distance over a candle or gas flame.
The cold metal cools the gases which rise from the flame,
and the water-vapour condenses to minute drops of moisture
on its surface, which in consequence immediately becomes
dim. The presence of carbon dioxide in the gases produced

28 INORGANIC CHEMISTRY

by the combustion can also be simply shown as follows : — If a
drop of lime water is taken up on the end of a glass rod, and
held some distance over the flame, it will at once become
milky, by the production in it of insoluble calcium carbonate
formed according to the equation —

Ca(0H)2 + CO2 = CaCOg* + HgO

Calcium hydroxide Carbon dioxide Calcium carbonate Water

Many elements when heated in the air to a sufficiently high
temperature take fire and burn, combining with the oxygen of
the air to form oxides. Thus sulphur, when heated, burns
with a characteristic blue flame to form the gas sulphur
dioxide,

S + O2 = SO2;

Sulphur Sulphur dioxide

and phosphorus burns with a brilliant white flame to form
clouds of phosphorus pentoxide P2O6, a solid which is pro-
duced according to the equation —

4P + 5O2 = 2P2O5

Phosphorus Phosphorus pentoxide

Even metals may burn in air if sufficiently heated. Thus
zinc can be made to take fire, combining with oxygen to form
the oxide ZnO according to the equation —

2Zn + 02 = 2ZnO,

Zinc Zinc oxide

the solid oxide appearing in the form of copious white fumes.
Iron wire, too, when raised to a very high temperature
burns readily with production of showers of sparks produced
by the vigorous union of the iron with oxygen to form the
oxide Fe203,

4Fe + 3O2 = 2Fe203

Iron Ferric oxide

The evolution of light does not necessarily accompany the
process of combustion, even when this is attended by con-

* The production of an insoluble substance {precipitate) may be con-
veniently represented in a chemical equation by underlining the formula of
the substance.

COMBUSTION 29

siderable production of heat. If a red-hot iron rod is allowed
to cool in the air, a scale will be found on the surface of the
rod where it has been raised to a high temperature. Although
no process of combustion is here evident, yet union has taken
place between the iron and the oxygen with production of the
same oxide as before. Again, if lead is melted in an iron vessel,
it will be found that the surface very soon becomes covered
with a solid scum. If this scum is raked off and a fresh bright
surface of lead exposed, a fresh scum makes its appearance,
which can in turn be removed. This solid scum is an oxide
of lead formed by the union of the metal with atmospheric
oxygen.

2Pb + O. = 2PbO

Lead Lead monoxide (litharge)

In the same way, when copper is moderately heated in air
it loses its metallic appearance and becomes covered with a
coating of oxide of copper by gradual union with the oxygen
of the air.

In animals a process of slow combustion constantly goes on,
which is in all essential respects closely related to the com-
bustion of fuel. The animal organism, like vegetable fuels,
consists very largely of substances containing carbon and hydro-
gen, which are slowly burned in the body by means of the
oxygen abstracted by the blood from the air, and conveyed by
it from the lungs to all parts of the body. The products
of this slow combustion in the body are carbon dioxide from
the carbon, and water from the hydrogen — that is, the
same substances as were produced by the rapid combustion
of ordinary fuel. These products are carried in the blood
back again to the lungs and leave the body in the expired
air. That they are contained in expired air in considerable
amount may easily be tested in the same way as before. If
we breathe gently on a cold bright object held in front of the
mouth, the surface will immediately become dim by the de-
position of moisture (which, however, is not all obtained from
the combustion of the tissues), and if we blow through a tube
the end of which is immersed in lime water, the lime water will
speedily become milky by the production of insoluble calcium
carbonate. Here the process of combustion goes on at the
comparatively low temperature of the animal body. Although

30 INORGANIC CHEMISTRY

no light is evolved, so much heat is given out in the process
that the temperature of the body is kept permanently above
the ordinary temperature of the atmosphere.

A given quantity of a given substance burning in air

will always give out the same amount of heat, no matter

whether it burns slowly or rapidly. This, of course, does not

imply that the same temperature will be reached in all cases.

If the combustion goes on slowly, the evolution of the fixed

amount of heat is spread over a long time, and the temperature

at any one instant is therefore not raised to a high degree on

account of the continual loss of heat by conduction and

radiation. On the other hand, if the combustion of the

substance takes place very rapidly, the same amount of heat

as before is produced in a much shorter time, and consequently

raises the reacting substances to a much higher temperature.

The temperature obtainable by the combustion of a given

amount of fuel depends, therefore, on whether the combustion

is made to take place slowly or rapidly, — in a large space,

where the cooling effect is great, or in a small space, where

the cooling effect is small. The combustion in the animal

body is a good example of a slow combustion spread over a

large space, the consequence being that the temperature of

the body is never raised to any high degree, although very

considerable quantities of material are burned.

•1

CHAPTER VI

FLAME

When iron burns in air or oxygen, much heat is given out in
the combustion, small incandescent particles being shot off in
the form of sparks. Similarly, when pure carbon is burned,
although the temperature is raised to
a very high degree, there is little or
no flame. On the other hand, when
coal-gas, or oil, or a candle is burned
in the air, the combustion is accom-
panied by the production of flame.

If we inquire into the nature of
flame, we find that what generally
goes under that name is really a
mixture of reacting gases raised to a
very high ttmperalure by the heat
generated in the action, so that unless
the reactmg substances are, during
combustion, in the gaseous state, the
combustion is not accompanied by pj^

the production of flame. When iron Diamona burning in CijgfTi
bums in oxygen, the oxygen is of

course a gas, but the iron is not, so that we have reaction
between a gas and a solid, which produces no flame. In the
same way, when charcoal burns in air or oxygen, although the
oxygen is gaseous, the carbon of which the charcoal almost
entirely consists remains solid at the temperature of the
combustion, and so no flame is produced.

In the case of coal-gas burning in the air, both the reacting
substances are gaseous, and the hot gases in the zone of
reaction form what we usually call a flame. If oil is burned
instead of coal gas, notwithstanding that oil is a liquid, the
combustion is attended by flame production. It must be

32 INORGANIC CHEMISTRY

remembered, however, that although oil is liquid at the ordinary
temperature, it is a gas at the temperature at which combustion
takes place. i

The oil, then, before reacting with the oxygen of the air, is
converted into a gas, and so the reaction in reality takes place
between gases, with consequent production of flame. Simi-
larly, although the wax of a candle is a solid, yet by the heat
of the reaction it is converted first into a liquid and then into
a gas, and a flame is therefore the result of the combustion.
Sulphur and phosphorus burn in air with production of flame
on account of the heat of the combustion of these solid sub-
stances melting them, and converting them into vapour before
they actually burn.

When fresh coal is put on a fire the combustion of the coal
is accompanied by the production of flame. Coal consists
principally of the element carbon as such, but contains besides
a considerable quantity of compounds of carbon with hydrogen,
which are generally called hydrocarbons. These compounds
are decomposed with production of combustible gases as the
coal is heated, and the gases thus produced react with the
oxygen of the air, the zone of reaction being marked by a
luminous flame. As the combustion proceeds the hydrocarbons
are practically all destroyed, and then there is nothing com-
bustible left behind but carbon, which, as we have seen, is not
volatile at the temperature of combustion, so that now there is
interaction between a solid and a gas, and the fire which is
what we call a bright or clear fire, simply glows without
production of flame.

Coal-gas burned in an ordinary burner gives a brightly
luminous flame. When burned in a bunsen burner, however,
which has the air holes open, it burns with a non-luminous
flame. If we shut off" the air supply at the bottom of the
bunsen, the luminosity reappears.

If we inquire into the cause of the luminosity of a coal-gas
flame burned at an ordinary burner, we find that it is pro-
duced by hydrocarbons which are present in the gas. At the
high temperature of the combustion, these hydrocarbons are par-
tially decomposed with separation of solid particles of carbon
within the flame, which, being heated to whiteness, render
the flame luminous.

When air is admitted through the air holes of a bunsen

FLAME

burner, we have seen that the luminosity of the flame disappears.

Therearetwopossiblereasonsforthis.
In the first place, the flame may be
cooled, because the coal-gas is diluted
with a considerable proportion of ni-
trogen from the air, which takes no
part in the chemical action. This
nitrogen must be heated along with
the other gases present; and since

O..

.'S

I.-.

Complete combuatlen

Partial combustion

No combustion

Fig. 5. — Bunsen Burner.

Gas is supplied through the tube G, and is
delivered rapidly through the small orifice ^,
air being at the same time sucked up through the
air-hole O. so that a mixture of gas and air passes
up the tube A. The collar C, which contains a
hole to correspond with O, may be made to admit
or shut off air by rotating it round A, In the
figure the air-hole is partially closed.

the heat given out by the
combustion of a given quan-

tity of coal-gas is always the Fig. 6.— Non-luminous Bunsen Flame.

same, the temperature pro- The letters A,g, and G have the same signifi-

At^r^^A ,^',U ««.«. U^ ««. UCrvU cance as in fig. 5. In the zone of no combus-

dUCed Will not be so high tj^n the mixfure of gas and air is below its

as when no nitrogen is preS- ignition temperature ; in the zone of partial com-

. J 1 /^ 1 bustion the coal-gas is not completely burnt ; in

ent. m tne SeCOna place, the zone of complete combustion there is excess

oxygen is SUPolied along *»^ ?*^' f°<^.i^« coal-gas U entirely burned to
. -^ o , , *^^ __- p carbon dioxide and water.

With the coal-gas. The coal-
gas has therefore a better chance of meeting the requisite quantity
of oxygen for its complete combustion (that is, the combus-

INORGANIC CHEMISTRY

tion of both carbon and hydrogen) when oxygen is supplied
within as well as without the flame, than when the oxygen supply
is derived entirely from the atmosphere outside the flame. In
the latter case, there is an insufficient supply of oxygen within
the flame for burning both the hydrogen and the carbon, so
that some of the carbon at first escapes
combustion, and is therefore seen in a solid
state at a white heat.

A luminous coal-gas flame may be made
non-luminous by mixing it with pure nitro-
gen, which plays no direct part in the com-
bustion at all, and in no way assists the
combustion of the coal-gas. Its action is

merely a cooling ac-
tion : the temperature
of the flame is kept so
low by the admixture
of nitrogen, that it
never reaches the
point necessary to de-
compose the illuminat-
ing hydrocarbons with
separation of carbon.
That cooling alone is
able to destroy the
luminosity in flame can be shown by hold-
ing a piece of fine copper wire gauze in a
slanting position, so that a feebly lumin-
ous bunsen flame plays against it. If

Fig. 7.
Candle Flame.

a. Area of no combustion-

b. Area of partial combus-

tion.

c. Area of complete com-

bustion.

Fig. 8.— Arrangement illustrating Principle of Bunsen Burner.

Gas is supplied through a mouth blow-pipe, the jet of which points directly
upwards through the tube of a glass filter-funnel. The lettering corresponds to
that^ in fig. 6. If G is rotated on its axis, the flame may be made to become
luminous. In this case, the jet of gas from f^ is broken against the side of the
funnel, and little or no air passes up A along with it.

the luminosity of the flame is not too great, it disappears^
entirely when the flame is thus brought into contact with the-

FLAME 35

wire gauze, which conducts away the heat very rapidly, and
thus lowers the temperature. That heat,* on the other hand,
confers luminosity on a non-luminous flame may be seen by
heating the tube of a bunsen burner to redness. Although,
before heating, the flame was non-luminous, after the tube
has been heated the flame becomes luminous. Here the
illuminating hydrocarbons have been decomposed with pro-
duction of particles of carbon, which in the flame are heated to
whiteness and consequently emit light.

In an ordinary gas burner, matters are so arranged that all
the carbon which separates is burned in the outer regions of
the flame, and no smoke is produced. If a large flame of coal-
gas is burned in a bunsen with the air supply completely shut
off, the flame will be seen to be smoky. Some of the carbon
particles have here escaped combustion altogether, and are
sent into the air as smoke, which may condense to what we
call soot.

The separation of solid particles is not the only possible
cause of luminosity in flame. Even though no particles
separate at all, a luminous flame may still be obtained by
sufficiently increasing the pressure on the reacting gases.
Thus hydrogen gas at the ordinary pressure of the atmosphere
burns with a non-luminous flame, the reacting substances and
all possible products of the reaction being gases at the tem-
perature of the flame ; but if the pressure of both gases is in-
creased to several atmospheres, the hydrogen flame becomes
luminous, notwithstanding the impossibility of the separation
of solid particles.

It has already been indicated that the temperature obtained
by the combustion of any substance depends upon the rate of
combustion, and on the size of the space in which the combus-
tion is eff*ected. Though the same quantity of gas is burned
in the same time, if the combustion is made in one case to
spread over a large area, and is in another case concentrated
to a very small space, the temperature of the flame in the
second case will be much higher than in the first. The same
amount of heat is produced in a much smaller space, and
therefore there is not the same loss of heat by conduction and
radiation. If we want, therefore, to get a hotter flame from
coal-gas than is given by complete combustion in the ordinary
bunsen burner, we can obtain it by means of a blow-pipe, in

INORGANIC CHEMISTRY

which there is an arrangement for rapidly supplying the air
necessary for the combustion of the gas. The coal-gas has
therefore not to spread out so far in order to meet the neces-
sary oxygen, for it is mixed at the nozzle of
the burner with all the oxygen it requires.
The reaction thus takes place in a smaller
space — that is, a smaller flame is produced,
but that flame is much hotter than if no
special air supply was given. To get a still
higher temperature from coal-gas, oxygen

Fig. 9. — Mouth Blow-pipe.

The mouth blow -pipe (shown in fig. 8) gives an
extra supply of air through the jet /, which may be
useful both for deflecting and concentrating the bunsen
flame from ^.

may be supplied in the blow-pipe instead
of air. The temperature rises above that
of the ordinary blow-pipe flame, because,
in the first place, no nitrogen has to be
heated, and, in the second place, the com-
bustion takes place in a much smaller
space, because the reacting gases are not
diluted with nitrogen. If we lower a
burning jet of coal-gas into a vessel filled
with pure oxygen, we can see that the
flame becomes much smaller than it was
in the air, and if it is tested it will be
found to be much hotter, because the
same amount of coal-gas is burned at the
jet in the two cases, and there is less loss of
case than in the first.

L J

■Oas

Air

Jb'ig. 10. — Blow-pipe.

Air is delivered rapidly
from a bellows or air-
blast through the narrow
central tube A^ coal-gas
passing up through the
space between this and
the outer tube B\ the
two only mix at the
mouth of the blow-pipe.
By this arrangement a
great amount of gas may
be burned in a small
s^ce, with consequent
high temperature.

heat in the second

FLAME

Fig. n.

When we consider that a flame consists of gases made
visible by their chemical interaction, we see that if these gases
were cooled below the temperature at which chemical action
between them is possible, the flame would disappear.

This can be shown readily by means of an ordinary non-
luminous bunsen flame and a piece of fine wire gauze. The
metal of the wire gauze exposes a large surface, and is a
good conductor of heat. If it is brought
into a flame the heat generated by the chemi-
cal action is conducted away along the wires
of the gauze and dissipated. If the mesh of
the gauze is fine enough, the gases are cooled
below the ignition point, and can, in conse-
quence, no longer interact.

Suppose the wire gauze to be brought
down horizontally on the flame (fig. ii).
It will be seen that the flame spreads slightly
beneath the wire gauze, but that it does not
penetrate to the upper surface of the gauze. That the mixture
of coal-gas and air which constitutes the unburnt gases from

the bunsen burner easily passes through the
holes in the gauze, although the flame does
not, may be proved by applying a light to
the upper surface of the gauze, when the
flame appears above the gauze as well as
below it. A combustible mixture must tfiere-
fore have passed through the gauze, although
the flame was unable to do so ; in other
words, below the gauze the gases were re-
acting chemically, but above the gauze they
were not.

If the wire gauze is held an inch or so
above the opening of an unlit bunsen burner,
and a light is applied aJjove the gauze (fig. 12),
the gases which pass through take fire, but the flame is not
transmitted downwards to the gases as they issue from the
mouth of the burner. The reason is the same as before :
the wire gauze lowers the temperature of the gases in its inter-
stices to a point below that at which they can react chemically.
The same cooling action of wire gauze may be shown in
another way. If burning alcohol is poured from a basin through

Fig. 12.

38 INORGANIC CHEMISTRY

the gauze, the part which passes through is extinguished,
although the alcohol above and on the gauze is blazing freely.

Should a bunsen flame play against the under surface of a
piece of wire gauze for a sufficient length of time to heat the
gauze to redness, it will be found that suddenly the flame passes
from beneath upwards through the gauze. This is due to the
gauze having lost its cooling power by itself becoming red-hot.
The principle of cooling reacting gases below their igni-
tion point by means of wire gauze is applied in miners' safety
lamps, which are usually named after their inventor, Davy.
The Davy lamp (fig. 13) consists of a small oil lamp over
which a cage or cylinder of wire gauze is
screwed down. This cage, which entirely
surrounds theflame, admits the free access
of air and free removal of the products of
combustion, but its cooling effect is such
that a flame cannot in general pass through
it. The reason why such safety lamps are
used in mines is that the mine sometimes
becomes filled with an inflammable mix-
ture of fire-damp and air. When mixed
with air in certain proportions this fire-
damp, or marsh gas, is not only inflam-
mable, but explosive. If a naked flame
were carried into such an explosive mix-
ture a very serious accident might happen.
If the flame, however, is guarded by a wire
Fig. 13.— Safety Lamp, gauze cage it is only the part of the com-
The iudu f fed from ihe bustible mixture inside the cage that is
«] MMrwjif d. burns wi.hin inflamed, and the flame produced is in-
Mi^e^'mr^st^'" "'' Capable of passing outwards through the
gauze.
The safety lamp is not an absolute protection against mine
explosions, for the following reason. If a shot is fired in the
neighbourhood of a safety lamp which is surrounded by an
explosive mixture of fire-damp and air, there is such a sudden
displacement of gas by the shot, that some of the burning gas
within the cage is driven bodily through the wire gauze with
such rapidity that the flame is not extinguished, and thus the
mixture outside the cage is fired. This may be shown with a
bunsen burner and a piece of wire gauze of somewhat coarse mesh.

FLAME

Gaf

If the gauze is brought down slowly on the flame, the flame will
not penetrate, and if we depress the gauze down to the actual
mouth of the burner the flame may be extinguished altogether.
If, on the other hand, we bring down the wire gauze rapidly on
the flame by a smart blow, the flame will be found to penetrate it.
In the latter case the reacting gases pass so rapidly through
the gauze that they have riot time to
be cooled by it below the ignition
point, and consequently they con-
tinue to react.

Since the combustion of coal-gas
in air is a chemical reaction which
takes place between the coal-gas and
the oxygen in the air, it is apparent
that we have been looking at this
chemical reaction from one side only.
Coal-gas and air will interact when
heated to the requisite temperature,
no matter in what way they are
mixed. For practical reasons we
usually have a jet of coal-gas issuing
into the surrounding air, but the
combustion would take place equally
well if a jet of air were made to issue
into an atmosphere of coal-gas.

This may be shown experiment-
ally by forcing a stream of air through
a small jet or burner and lowering
this into a jar filled with coal-gas,
having previously lit the jet of air at the flame of the coal-gas
burning in air at the top of the jar (fig. 14).

In a similar way we very frequently, on account of our every-
day practice in conducting a chemical action, are one-sided in
our terminology. Thus we say that coal-gas is a combustible
substance and that air is a supporter of combustion ; but as we
have just seen, the terms might with equal propriety be
reversed. When we say that a substance is combustible^ with-
out further qualification, we imply that it will burn in air or
oxygen ; and when we say that a gas is a supporter of combus-
tion^ we mean that it will react chemically with the same
substances, and in the same way, as air does.

vi^y

Fig. 14. — Combustion of Air
in Coal-gas.

A. Flame of air burning in coaI«

gas.

B. Flame of coal-gas burning in

air.

CHAPTER VII

NEUTBALISATION

In carbonic acid and calcium hydroxide we have met with
examples of two very important classes of chemical substances
— namely, acids and bases. They are, however, scarcely to be
called typical examples, and in the present chapter we shall
be concerned with the soluble acids and bases which are more
commonly in use in the laboratory, and possess better defined
properties. The common soluble acids are —

Chemical Name

Formula

Common Name

Sulphuric acid

H2SO4

Oil of vitriol

Nitric acid

HNOs

Aquafortis

Hydrochloric acid

HCl

Muriatic acid

The common soluble bases or alkalies are —

Sodium hydroxide NaOH Caustic soda

Potassium hydroxide KOH Caustic potash

Ammonium hydroxide NH4OH Ammonia

A convenient test for soluble acids and alkalies is their
action on a solution of litmus, a purple colouring-matter ex-
tracted from certain lichens. Acids turn the purple litmus
solution red, alkalies turn it blue. Papers impregnated with
litmus are extensively used instead of litmus solution in testing
for acids and alkalies, and are known on this account as test-
papers. Acids and alkalies can be distinguished from each
other not only by their action on litmus, but also by their
taste. Acids have a sour taste; alkalies have a somewhat
soapy flavour.

A curious characteristic of acids and alkalies is the power
they possess of destroying or neutralising each other's pro-
perties. If we mix a solution of any of the above acids with a

NEUTRALISATION 41

solution of any of the above alkalies in the proper proportions,
the resulting solution will neither redden litmus nor turn it
blue, and will be neither sour nor alkaline to the taste.
The characteristic properties of the original substances will
have vanished, and new properties will have made their appear-
ance. The solution obtained by the mixing of the acid and
alkaline solutions leaves purple litmus unaffected, and possesses
a taste resembling that of salt. The reciprocal action of acid
and base is called neutralisation, and is a very important
chemical process.

When we inquire more closely into the phenomena of
neutralisation, we find that the definite proportions, which we
have already seen to exist in the processes of burning and
slaking lime, are equally evident in the reactions between
acids and alkalies. For example, if we take two solutions
of acids (say one of hydrochloric acid and the other of
sulphuric acid) and find that three times as much of the
sulphuric acid is required to neutralise a given weight of caustic
soda as is required of the hydrochloric acid solution ; then
this same relation will be found with regard to any quantity
of any alkali. That is, we may substitute caustic potash or
ammonia for the caustic soda, and taking any quantity of
them we please, we shall always find that three times as
much of the given sulphuric acid solution is required to
neutralise the alkali taken as is required of the given hydro-
chloric acid solution. The quantities of different acids
which are capable of neutralising the same quantity of a
given alkali are said to be equivalent to each other : in the
above instance the quantity of hydrochloric acid in one
volume of the hydrochloric acid solution is equivalent to
the quantity of sulphuric acid in three volumes of the sul-
phuric acid solution. Using this conception of equivalence,
we may say then that quantities of different acids which are
equivalent with respect to one alkali are equivalent with
respect to all alkalies.

Similarly, we find that quantities of different alkalies which
are equivalent with respect to a given quantity of a given
acid will neutralise the same quantity of any acid.

By an extension of the term equivalent we say that quan-
tities of acid and alkali which are capable of neutralising each
other are equivalent to each other.

INORGANIC CHEMISTRY

The substances formed by the mutual neutralisation of an
acid and an alkali are water and a salt. The salts, of which
ordinary salt is a typical example, form another very impor-
tant class of chemical substances. They are solids under
ordinary circumstances, generally possess a salt taste, and
yield solutions which are mostly neutral to litmus.

As examples of salt-formation we may take the neutralisa-
tion of caustic soda by the three acids mentioned above,
and express the reactions by means of equations.

NaOH + HCl

Sodium hydroxide Hydrochloric acid
40 36.5

NaOH -h

Sodium hydroxide
40

HNO3

Nitric acid

2NaOH + H2SO4

Sodium hydroxide Sulphuric acid
80 98

NaCl +
Sodium chloride

58-5

NaNOs +

Sodium nitrate
85

Na2S04

Sodium sulphate
142

HjjO

Water
18

H2O

Water
18

+ 2H2O

Water

Here the alkali is the same throughout, and the acid is
varied. We may now write a similar set of equations for the
neutralisation of hydrochloric acid by the various alkalies.

HCl + NaOH =

Hydrochloric acid Sodium hydroxide
36.5 40

NaCl +

Sodium chloride
58.5

HCl + KOH = KCl +

Hydrochloric acid Potassium hydroxide Potassium chloride
36.5 56 74.5

HCl 4-

Hydrochloric acid

36.S

NH4OH

Ammonium
hydroxide

= NH4C1 +

Ammonium chloride

53-5

H2O

Water
18

H2O

Water
18

H3O

Water
18

By taking the other pairs of acid-alkali we get still other
salts, the equations representing the formation of which the
student should write for himself. A list of all the salts formed

NEUTRALISATION

by the neutralisation of the common acids and alkalies is given
below.

Chemical Name

Formula

Common Name

Sodium chloride

NaCl

Salt

Potassium chloride

KCl

Chloride of potash

Ammonium chloride

NH4CI

Sal-ammoniac

Sodium nitrate

NaNOg

Chili saltpetre

Potassium nitrate

KNO3

Saltpetre or nitre

Ammonium nitrate

NH4NO3

Nitrate of ammonia

Sodium sulphate

Na2S04

Salt-cake

Potassium sulphate

K2SO4

Sulphate of potash

Ammonium sulphate

(NH4)2S04

Sulphate of ammonia

All of these salts are solids which are soluble in water, are
neutral to litmus, and have a taste resembling that of ordinary
salt.

If we now consider the quantities of substances taking
part in the various actions, we see that to neutralise 40 parts
of caustic soda are required —

36.5 parts of hydrochloric acid
63 „ nitric acid
49 „ sulphuric acid.

These quantities are therefore equivalent.

Again, in order to neutralise 36.5 parts of hydrochloric acid
we require —

40 parts of sodium hydroxide
56 „ potassium hydroxide
35 „ ammonium hydroxide.

These quantities of the various alkalies are therefore equi-
valent to each other. Not only, however, are these quan-
tities of the acids and alkalies equivalent as compared each
with a substance of like kind — i.e, acid with acid, and alkali
with alkali, — they are also equivalent when alkali is compared
against acid. Thus 35 parts of ammonium hydroxide are
equivalent to 63 parts of nitric acid — Le, these quantities will
exactly neutralise each other; and 56 parts of potassium

44 INORGANIC CHEMISTRY

hydroxide are equivalent to 49 parts of sulphuric acid, as
reference to the equations will show.

Using formulae instead of the formula-weights connected
with them, we see that NaOH, KOH, and NH4OH are
equivalent to one another. HCl and HNO3 are also equi-
valent, not only to each other but also to NaOH, etc. It
is different in the case of sulphuric acid. The formula
H2SO4 represents 98 parts of sulphuric acid, which the equa-
tions show will neutralise twice as much caustic soda as 36.5
parts of hydrochloric acid. The formula H2SO4 is therefore
equivalent to 2 HCl, and also to 2HNO8. When compared
against alkalies, the quantity expressed by H2SO4 neutralises
twice as much as the quantities expressed by the formulae
KOH, NaOH, and NH4OH; H 2 SO 4 is therefore equivalent
to 2NaOH, 2KOH, and 2NH4OH.

CHAPTER VIII

THE COMMON ACIDS AND BASES

The three common acids — sulphuric acid, nitric acid, and
hydrochloric acid — are usually called the strong mineral acids,
to distinguish them from other acids of a weaker nature derived
from minerals, and from the vegetable acids.

Sulphuric Acid, H2SO4. — This is the most important of all
acids, and is prepared commercially on an enormous scale. It
is not only the commonest acid, but the cheapest, and
nearly all other acids are prepared from it by its action on
certain salts. The crude commercial acid is known as oil of
vitriol, and forms an oily liquid which is usually somewhat
brown in colour.

This oil of vitriol contains, besides other impurities, about
20 per cent, of water. Pure sulphuric acid is a heavy, perfectly
colourless liquid. It is a powerfully corrosive substance, and
destroys all animal and vegetable substances with which it
comes in contact, blackening and charring them by partially
converting them into charcoal. This charring may easily be
observed by bringing sulphuric acid into contact with sugar,
papei:, or wood.

If it is poured into water, a great amount of heat is evolved,
the rise of temperature being sufficient to convert a portion
of the water into steam. In diluting it with water, care must
therefore be exercised, as the production of steam may be so
sudden as to cause portions of the corrosive liquid to fly about.
The undiluted acid is often spoken of as strong sulphuric acid,
and the diluted acid as weak sulphuric acid. It is better,
however, to reserve the terms strong and weak to indicate the
inherent strength or weakness of pure acids, using the terms
concentrated and dilute to express the degree with which the
pure substances are mixed with water.

INORGANIC CHEMISTRY

Pure sulphuric acid boils at a temperature above 300°,
on which account it can be used
to drive out more volatile acids
from their salts, as will be seen
below.

Nitric Acid, HHO3. — This
acid is produced by the action
of concentrated sulphuric acid on
sodium nitrate. In the labora-
tory the two substances may be
heated together in a retort. The
i*urK Mid is htated in the giais r«wt sodlum nitrate dissolves in the

Vf, from which ihe mmt acid dmils off, , , ■, , .. .

being condcnsea la ihc liquid foim in sulphuHC acid on heatmg, and
ajet'^mwr'p'ia'ylng^in'i^*'" "'** '"'' ^^^ nitric acid produced by their
interaction boils off, the vapour
being liquefied in a cooled receiver. The equation for the
action is —

H,SO, + NaNOa = NaHSO^ + HNOg

Sulphuric acid Sodium nitrate Sodium hydrogen sulphate Nitric acid
98 85 120 63

Nitric acid boils at a temperature below that of boiling water,
whilst all the other substances concerned in the reaction boil

entlieiy lorrounded liy brick- work ,
diiiila off iiliqueliEd by pauing Ihi

). — Commercial I'reparalion of Niltic Acid.

' is heated liT the flames ind Ct nasct frin
iff [liroagh the flue ^. The nitric acid i

at a much higher temperature ; so that the t
off and is thus separated from them.

ITHE COMMON ACIDS AND BASES

Nitric acid produced in this way has always a reddish-yellow
colour, which is due to the decomposition of a little of the acid
by heat during the distillation. When pure it is a colourless
liquid, not so heavy as sulphuric acid, although much heavier
than water, with which it mixes with considerable evolution
of heat. The concentrated acid is powerfully corrosive, attack-
ing both animal and vegetable substances, usually staining
them bright yellow.

HydrocMoric Acid, HCl. — This acid is formed by the action
of sulphuric acid on common salt (sodium chloride), the equa-
tion being —

H2SO4 + 2NaCl « Na2S04 + 2HCI

Sulphuric acid Sodium chloride Sodium sulphate Hydrochloric acid
98 117 142 73

It is a gas, so that when produced, it
escapes from the other substances con-
cerned, which are solid or liquid. As
gases are extremely bulky and incon-
venient to work with, pure hydrochloric
acid is scarcely ever employed. The gas
is led off as it is formed into water,
which can absorb many hundred times
its own volume. The solution of the gas
in water is accompanied by a consider-
able rise in temperature. What is known
as concentrated hydrochloric acid in the
laboratory is an aqueous solution con-
taining about 35 per cent, of the pure acid.
Dilute hydrochloric acid contains not more
than 20 per cent. The gas and the solution
derived from it are without colour when
pure, and are not nearly so corrosive in
their action as sulphuric or nitric acids.

Sodium hydroxide, NaOH, and Pota.8-
sium hydroxide, EOH, are the common
fixed — that is, non - volatile — alkalies.
They are white solids which are freely
soluble in water. Their concentrated
solutions are strongly caustic and cor-
rosive, their dilute solutions less so.

Fig. 17. — Preparation
of Gaseous Hydro-
chloric Acid.

Salt is introduced into
the flask Ff and concen*
trated sulphuric acid poured
through the thistle funnel
7", the tube of which nearly
reaches the bottom of the
flask. The hydrochloric
acid produced escapes
through the exit tube £.
Heat may be applied if
necessary by means of a
burner placed beneath the
flask.

48 INORGANIC CHEMISTRY

Ammonium hydroxide, NH4OH, only exists in aqueous
solutions, which are prepared by dissolving ammonia gas in
water, the equation for the action being —

NH3 + H2O = NH4OH

Ammonia Water Ammonium hydroxide

It was formerly called the volatile alkali on account of the ease
with which it breaks up again into ammonia gas and water,
according to an equation which is the reverse of that just
given. The solution of ammonia may be easily distinguished
from the solutions of the other alkalies by its characteristic smell,
sodium hydroxide and potassium hydroxide being practically
inodorous.

These alkalies, as has already been mentioned, form a
special subdivision of a much larger class of substances called
bases. All bases are hydroxides^ and have the power of react-
ing chemically with acids to form salts and water, just as the
alkalies do. The other bases differ from the alkalies in not
being freely soluble in water. Calcium hydroxide, for example,
although not usually termed an alkali, is, like sodium hydroxide,
a base, and capable of turning litmus blue and of neutralising
acids. Thus, when calcium hydroxide is treated with hydro-
chloric or sulphuric acids, neutralisation takes place according
to the following equations : —

Ca(0H)2 + 2HCI = CaCla + 2H2O

Calcium hydroxide Hydrochloric acid Calcium chloride Water

Ca(0H)2 + H2SO4 = CaS04 + 2H2O

Calcium hydroxide Sulphuric acid Calcium sulphate Water

Calcium hydroxide, however, is only very sparingly soluble in
water, and although its solution has an alkaline reaction, is not
termed an alkali, but an alkaline earth. Strontium hydroxide
and barium hydroxide Ba(0H)2 are similar to calcium
hydroxide in this respect, and are classed with it as alkaline
earths. Owing to this classification of their hydroxides, the
metals sodium and potassium are often referred to as metals
of the alkalies, or alkali metals; and the metals calcium,
strontium, and barium, as metals of the alkaline earths.

Some bases, such as zinc hydroxide Zn(0H)2, possess no
alkaline reaction, and are altogether insoluble in water. They

THE COMMON ACIDS AND BASES 49

nevertheless react with acids to form salts in precisely the
same way as the soluble bases. Thus zinc hydroxide and
hydrochloric acid at once give zinc chloride and water:

Zn(0H)2 + 2HCI = ZnCla + 2H2O

Zinc hydroxide Hydrochloric acid Zinc chloride Water

and similarly,

Zn(0H)2 + H2SO4 = ZnSO^ + 2H0O

Zinc hydroxide Sulphuric acid Zinc sulphate Water

CHAPTER IX

•>«ir

The process of ordinary combustion is a process of oxidation
— that is, a process of union with oxygen ; and if elements are
the substances which are burned, the compounds produced are
called oxides. Thus the element carbon burns to form carbon
dioxide, hydrogen to form hydrogen oxide or water, zinc to form
zinc oxide, iron to form an oxide of iron, phosphorus to form an
oxide of phosphorus, and sulphur to form an oxide of sulphur.
Some elements form more than one oxide. Carbon, when
burned in air or in oxygen, bums for the most part directly
to carbon dioxide; but in some cases where there is not a
sufficient supply of oxygen, another oxide, carbon monoxide,
may be produced according to the equation —

2C + O2 = 2CO

Carbon monoxide

This monoxide differs altogether in its properties from the
dioxide. The dioxide will not itself burn in^'air, and at once
extinguishes a taper. The monoxide, although it extinguishes
a taper, burns in air with a blue flame when a light is applied
to it. The dioxide, as we have seen, turns lime-water milky :
the monoxide has no effect on lime-water.
When we compare the equations

C + O2 = CO2

Carbon dioxide

2C + O2 = 2CO

Carbon monoxide

wc see that for a given amount of carbon, twice as much
oxygen is required to produce the dioxide as to produce the
monoxide ; and it is only when oxygen is present in insufficient
quantity that the monoxide seems to be produced.

The production of carbon monoxide can easily be effected
by passing a current of the dioxide over red-hot carbon. The

OXIDES 51

carbon and the carbon dioxide react with each other according
to the equation —

C + CO2 = 2CO

Carbon Carbon dioxide Carbon monoxide

This action may be observed in the burning of a clear coal
fire. The glowing coals consist almost entirely of carbon;
and carbon is a substance which burns without flame. It
will be seen, however, that little blue flames flicker on the
top of the fire. These blue flames are due to the combustion
of carbon monoxide. When air enters the fire at the bottom
of the grate, its oxygen combines with the carbon of the coal
to produce carbon dioxide, which, rising higher in the fire,
comes into contact with red-hot carbon, and is converted into
carbon monoxide. The carbon monoxide continues to ascend
until it escapes at the top of the fire, where it mixes with more
air, and once more produces carbon dioxide.

2CO +02= 2CO2

Carbon monoxide Oxygen Carbon dioxide

The formulae of the oxides of carbon indicate that the monoxide
for a given quantity of carbon contains less oxygen than the
dioxide. We therefore say that the carbon in carbon monoxide
is at a lower stage of oxidation than the carbon in carbon
dioxide, or that carbon in the monoxide is less highly oxidised
than carbon in the dioxide.

Another example of an element which combines with oxy-
gen in more than one proportion is to be found in sulphur.
When sulphur burns in air or in oxygen under ordinary
conditions, the chief product of the combustion is sulphur
dioxide SO 2, but at the same time a small quantity of another
oxide, sulphur trioxide SO 3, is formed. But even though
there is a large excess of oxygen, very little of the oxide SO 3 is
produced, by far the greater part of the sulphur remaining
in the less highly oxidised stage of SO 2. In this respect
sulphur differs altogether from carbon. The tendency for
carbon is to pass on combustion into the higher stage of
oxidation, while the tendency for sulphur is to remain at the
lower stage of oxidation.

When an element unites with oxygen in more than one
proportion, the various oxides produced are frequently dis-
tinguished from one another by means of the Greek numerals

INORGANIC CHEMISTRY

prefixed to the word oxide. The numbers indicate the
numbers of oxygen symbols contained in the formula, thus :

Carbon Monoxide

Carbon Dioxide

CO2

Sulphur Dioxide

SO2

Sulphur Trioxide

. SOs

Nitrogen Tetroxide .

. N,0,

Nitrogen Pentoxide .

. N,0,

Phosphorus Pentoxide

. P2O,

Although many elements combine directly with oxygen
when raised to a certain temperature, this is by no means
invariably the case. The metal mercury when kept at its
boiling point in presence of air or oxygen, slowly unites with
the oxygen to form a quantity of mercuric oxide.

2Hg

2HgO

Mercuric oxide

Ifnve heat this oxide, however, to a slightly higher temperature
than that^at which it was formed, it breaks up again into
metallic mercury and oxygen, which cannot be made to unite
by further raising the temperature. This decomposition of
mercuric oxide is sometimes used as a means of preparing a
small quantity of oxygen (p. 127).

If instead of the metal mercury we take the metal silver
and heat it, we find that at no temperature will it combine
with oxygen to form oxide of silver. The oxide of silver may
be formed, however, indirectly; but when heated, it decom-
poses into metallic silver and oxygen.

2AgoO = 4Ag
Silver oxide Silver

+ O3
Oxygen

The oxide ZnO, on the other hand, formed by the com-
bustion of zinc, cannot be decomposed by heat, no matter how
high the temperature is raised.

Zinc is therefore an example of a metal which forms an
oxide by direct combustion which is perfectly stable towards
heat Mercury is a metal which slowly oxidises at a moderate
temperature to form an oxide, but this oxide easily decom-
poses again when the temperature is raised. Silver is an
example of a metal which will not form an oxide at all by

OXIDES S3

direct union with oxygen, and is reproduced from the oxide
by heat when this has been formed indirectly.

We have already had instances of the production of oxides
by decomposition. We have seen that both calcium carbonate
and calcium hydroxide decompose at a red heat with formation
of calcium oxide, carbon dioxide, and water according to the
equations —

CaCOg = CaO + CO 2

Calcium oxide Carbon dioxide

Ca(0H)2 = CaO + HgO

Calcium oxide Water

Many other carbonates and hydroxides decompose in a similar
manner, giving off respectively carbon dioxide and water as
gases, and leaving behind a residue of a metallic oxide.

Oxides may sometimes be formed by the careful decomposi-
tion of nitrates by heat. Lead nitrate when heated gives off
gaseous nitrogen peroxide and oxygen, and leaves a residue
of lead monoxide (litharge), the decomposition taking place
according to the equation —

2Pb(N03)2 = 2PbO + 4NO2 + O2

Lead nitrate Lead monoxide Nitrogen peroxide Oxygen

Similarly, mercuric nitrate decomposes when heated to form
mercuric oxide, the same gases being given off. In this case
the heating must be done very gently, otherwise the tempera-
ture would rise to such a degree as to decompose the mercuric
oxide which is formed.

The oxides of the elements form an extremely important
class of substances, which may be divided into several sub-
classes, the chief of which are the basic oxides, and the acidic
oxides. The basic oxides produce bases by union with water,
or can be derived from bases by depriving them of water.
Acidic oxides, on the other hand, produce acids by union with
water, or can be formed by depriving acids of water.

An example of a basic oxide has already been met with in
calcium oxide. This substance unites readily with water to
form the base calcium hydroxide, and can be reproduced from
calcium hydroxide by heating that substance to a red heat, the
reversible equation being

CaO + H2O ^ Ca(0H)2

54 INORGANIC CHEMISTRY

Zinc oxide is another example of a basic oxide. It, however,
cannot be made to unite directly with water, to give the base
zinc hydroxide Zn(0H)2, but it can easily be derived from
zinc hydroxide by heating, the equation being —

Zn(0H)2 = ZnO + HgO

Zinc hydroxide Zinc oxide

Examples of acidic oxides, or acid anliydrides, as they are
sometimes called, have been seen in carbon dioxide and
sulphur dioxide, which unite with water to form carbonic acid
and sulphurous acid respectively, according to the equations —

CO2 + H2O ^ H2CO3

Carbon dioxide Carbonic acid

(Carbonic anhydride)

SO2 + HoO ^ HgSOs

Sulphur dioxide Sulphurous acid

(Sulphurous anhydride)

The acids formed from these anhydrides cannot, however, be
obtained in the pure state, on account of the ease with which
they decompose into the anhydrides and water, the equations
for the decomposition being the reverse of those for their
formation. Other acid anhydrides have a much greater
tendency to remain combined with water — e,g, sulphuric
anhydride SOg, which unites with water as follows: —

SO3 + HjjO = H2SO4

the sulphuric acid so produced only regenerating the anhydride
with difficulty.

Some oxides, such as nitric oxide NO, are said to be neutral
oxides — that is, neither acidic nor basic ; but such oxides are
exceptional, nearly all oxides possessing a more or less strongly
marked acidic or basic character.

The chemical distinction between metals and non-metallic
elements is chiefly based on the character of their oxides.
The non-mcta/s form no basic oxides ; every metaJ^ on tfu other
handy forms at /east one basic oxide (p. 220).

CHAPTER X

FORMATION AND DECOMPOSITION OF SALTS

We have seen in a previous chapter how salts may be pro-
duced by the mutual neutralisation of acids and bases. This
is not the only way in which salts may be formed. Bases,
which are metallic hydroxides, are related to metallic oxides
by having in addition water in their composition. If now,
instead of the base itself, we bring a basic oxide in contact
with an acid, the two substances interact to produce a salt, the
only difference from the ordinary production of salts by
neutralisation being that less water is formed in the action.
Thus zinc oxide is at once acted on by hydrochloric acid, and
by sulphuric acid, the reactions taking place according to the
following equations : —

ZnO + 2HCI = ZnClg + H2O
ZnO + H2SO4 = ZnS04 + H^O

Zinc chloride and zinc sulphate are the salts produced. If
we compare these equations with the following, which show
the production of the same salts from zinc hydroxide : —

Zn(0H)2 + 2HCI = ZnClg + 2H2O
Zn(0H)2 + H2SO4 = ZnSO^ + 2H2O

we see that twice as much water is produced in the second
case as in the first ; but that the other products of the reaction
are exactly the same.

All basic oxides are acted on by acids in this way, and
produce the same salts as would be produced from the
hydroxides which are derived from these basic oxides by the
addition of water.

Another important method of salt production is by the
aetion of acids on metals. Acids are frequently said to dis-
solve metals, but it must be noted that the metal is not
dissolved by the acid in the same sense as salt is dissolved

56 INORGANIC CHEMISTRY

by water. If we evaporate a solution of salt to dryness, the
original salt is obtained as a residue. If we evaporate a
solution of zinc in sulphuric acid to dryness, we do not obtain
the metal zinc, but the salt zinc sulphate. A metal when
it dissolves in acid is altogether changed ; it is converted by
the acid into a soluble salt of the metal, which then dissolves —
i,e. the solution obtained by acting on an acid with a metal is
a solution of a salt, and not a solution of the metal itself. The
different acids vary in their action on metals. Hydrochloric
acid and dilute sulphuric acid act in general least readily;
nitric acid, and concentrated sulphuric aqid, most readily.
Heat in every case promotes the action of an acid on a
metal.

Zinc is attacked by all three acids, whether dilute or con-
centrated. With dilute sulphuric acid the action is

Zn + H2SO4 = ZnS04 + Hg

Zinc Zinc sulphate Hydrogen

With hot concentrated sulphuric acid the action is

Zn + 2H2SO4 = ZnS04 + SO2 + 2H2O

Zinc Zinc sulphate Sulphur dioxide Water

When concentrated sulphuric acid acts on a metal, sulphur
dioxide and water are produced at the same time as a sulphate
of the metal, instead of the hydrogen gas which appears when
dilute sulphuric acid is employed. Sulphur dioxide is a
non-combustible gas which is moderately soluble in water,
and can easily be recognised by its characteristic smell of
burning sulphur. Hydrogen is a combustible gas, which,
when pure, is free from smell. It will be noted that twice
as much sulphuric acid is required to convert zinc into zinc
sulphate when the acid is concentrated, as when the acid
is dilute.

With hydrochloric acid, whether concentrated or dilute,
the action is

Zn + 2HCI = ZnCla + Hg

Zinc Zinc chloride Hydrogen

This type of reaction is always met with when hydrochloric
acid acts on a metal. The chloride of the metal and
hydrogen gas are the only products.

FORMATION AND DECOMPOSITION OF SALTS 57

The action of nitric acid on metals is usually somewhat
complicated, a nitrate of the metal being almost always formed
together with water and some compounds of nitrogen (pp. 104-7).
The fact which the student should specially note is that hydrogen
is not evolved when a metal dissolves in nitric acid.

It has been stated above that hydrochloric and dilute
sulphuric acid are usually somewhat less active, so far as
the solution of metals is concerned, than nitric or concent-
rated sulphuric acid. The metals copper, mercury, and
silver, for instance, are not attacked by dilute sulphuric or
by hydrochloric acid, but will readily dissolve in nitric or
in hot concentrated sulphuric acid, with formation of nitrates
or sulphates. This difference in action will be referred to
again in the sequel.

There are many other ways of producing salts; but those
just given are the most important, and will suffice at present
for the student.

Decomposition of salts by heat. — We have seen that the
metallic salts may be derived on the one hand from an acid,
and, on the other, from the hydroxide or basic oxide of a metal,
or even from the metal itself. When such metallic salts are
heated, they very frequently decompose with separation of
the metallic or basic part from the acidic part. We have
already had an instance of this in the action of heat on calcium
carbonate, which decomposes into the metallic oxide CaO and
the acidic anhydride CO 2, according to the equation —

CaCOg = CaO + CO2

Practically all carbonates decompose in this manner at a red
heat or below it, the chief exceptions being sodium carbonate
Nag CO 3, potassium carbonate Kg CO 3, and barium carbonate
BaCOg. Again, we have seen that mercuric nitrate, when
heated, decomposes with production of mercuric oxide HgO,
and gases originally derived from the nitric acid.

2Hg(N03)2 = 2HgO + 4NO2 + O2

Mercuric oxide Nitrogen peroxide Oxygen

Nearly all metallic nitrates decompose in a similar way, leaving
behind a residue of a metallic oxide, the chief exceptions being

58 INORGANIC CHEMISTRY

the nitrates of potassium and sodium KNO3 and NaNOa
(see p. 163).

The chlorides derived from hydrochloric acid contain no
oxygen, and can therefore, when heated alone, leave no
residue of a metallic oxide, although sometimes, when heated
in air, they are partially converted into oxides by the action
of the oxygen of the air. The metallic chlorides very fre-
quently withstand a high degree of heat before any change
occurs, and then they often merely vaporise without de-
composition.

The siQpliates of the metals usually withstand a consider-
able amount of heating without decomposition, but at a dull
red heat they mostly decompose, giving off an oxide of
sulphur, and leaving behind a residue of metallic oxide. The
sulphates of potassium, sodium, calcium, barium, and lead may,
however, be heated to redness without decomposition.

CHAPTER XI
POSITIVE AND NEGATIVE BADIOALS

In order to deal in a simple way with many actions in which
acids, bases, and salts take part, it is convenient to conceive
these substances to be composed each of two parts — namely,
a positive radical and a negative radical. By radical we
mean an element or group of elements in definite proportions
which occurs in a number of similar compounds ; thus all the
sodium salts contain the simple radical Na, all the calcium salts
contain the simple radical Ca, and all the ammonium salts
contain the compound radical NH^. Similarly, all chlorides
contain the simple chloride radical CI, and all nitrates the com-
pound radical NO 3. Positive radicals are so called because
they can under certain circumstances unite with and carry
charges of positive electricity. Negative radicals in a similar
manner can unite with and carry charges of negative electricity.

Table of Common Radicals and Ions

Positive Radicals. Negative Radicals.

(Kations.) Monad (Anions.)

H- Hydrogen OH' Hydroxide

Na* Sodium

K* Potassium

N H ^ Ammonium

Ag* Silver

Hg' Mercurous

Hg" Mercuric

Ca" Calcium

Ba** Barium

Zn" Zinc

Cu" Copper

Pb" Lead

Sn" Stannous

Fe" Ferrous

Fe"* Ferric

Al"' Aluminium

Dyad

Triad

CI'

Chloride ^

Br'

Bromide

Iodide

CIO,'

Chlorate

NO3'

Nitrate

HS'

Hydrosulphide

SO/

Sulphate \

S«03"

Thiosulphate ^

SO3"

Sulphite I

Sulphide \

[O"

Oxide] I

CO,"

Carbonate

PO/"

Orthophosphate

6o INORGANIC CHEMISTRY

The table on preceding page gives a list of the common posi-
tive and negative radicals. The symbols of the radicals are to be
interpreted in the usual way; thus Na stands for 23 parts by
weight of sodium; CI stands for 35.5 parts by weight of
chlorine; NH4 stands for 14 parts of nitrogen combined with
4 parts of hydrogen. The dots and dashes attached to the
symbols indicate the amounts of electricity with which the
quantities expressed by the symbols may be associated. Each
dot stands for unit charge of positive electricity, each dash
stands for an equal charge of negative electricity.

The symbol Ca", then, indicates that, as compared with the
symbol Na*, 40 parts of the calcium radical are charged with
twice as much positive electricity as 23 parts of the sodium
radical. Again, comparing the symbols Na* and CI', we see
that 23 parts of the sodium radical are charged with just as
much positive electricity as would neutralise the negative
electricity carried by 35.5 parts of the chloride radical.

A consideration of the electrical charges associated with
each radical gives us some notion of how these radicals might
be expected to combine with each other. It is a fundamental
electrical fact that charges of the same kind repel each other —
that is, positive repels positive, and negative repels negative;
whilst charges of opposite kinds attract each other — that is,
positive attracts negative. We should therefore expect to find
that radicals charged with positive electricity combine, not
with other radicals charged with positive electricity, but with
radicals charged with negative electricity; and similarly, that
radicals charged with negative electricity combine, not with
other negatively charged radicals, but with the oppositely
charged positive radicals. This indication is in accordance
with fact. Acids, bases, and salts are all composed of posi-
tive radicals combined with negative radicals.

We know that acids, bases, and salts are under ordinary
circumstances electrically neutral — that is, they have, on the
whole, neither a charge of positive electricity nor a charge
of negative electricity. The charges of positive and negative
electricity, therefore, which they are capable of assuming, must
be such as will exactly neutralise each other. Now, we have
expressed the unit charge by a dot if the electricity is posi-
tive, and by a dash if the electricity is negative. In the
formulae each dot is thus capable of exactly neutralising each

POSITIVE AND NEGATIVE RADICALS 6i

dash. If the substance expressed by a combination of symbols,
then, is to be exactly neutral, the formulae must contain equal
numbers of dots and dashes.

This equivalence in the number of dots and dashes con-
tained in the formula for an acid, base, or salt gives us an easy
means of writing chemically correct formulae for these sub-
stances. Thus, if we are asked what is the formula of a salt
containing the sodium radical and the chloride radical, we can
say at once that it must be Na*Cr, because each sodium radical
has one positive charge, and each chloride radical one negative
charge, so that they must be associated in equal numbers in
order to make an electrically neutral compound. Again, if we
are required to give the formula of a salt containing calcium as
the positive radical, and chloride as the negative radical, we
must write Ca**Cl2'. In order to balance the two positive
charges on the calcium radical, we must have two chloride
radicals each associated with one negative charge. Other
examples of a similar nature will be given in the course of
the chapter.

Salts receive names which indicate the positive and negative
radicals they contain. Thus the chemical name of common
salt Na'Cr is sodium chloride^ the chemical name of salt-
petre or nitre K'NOa', '^^ potassium nitrate^ the chemical name
for calc-spar, etc., Ca^COg" is calcium carbonate, and so on. In
naming the salt the positive radical is placed first and the
negative radical second.

Adopting this system, we are enabled at once to write the
correct formula for a salt as soon as its systematic chemical
name is given. We are required, for example, to write the
formula of zinc nitrate. We know that the positive radical of
this salt is zinc Zn", and that the negative radical is nitrate NO 3'.
In order that the number of dashes on the negative portion
shall be equal to the two dots on the zinc radical, we must
take two of the nitrate radicals, and so we write the formula
of zinc nitrate as follows: — Zn**(N03')2. Again, we may be
asked to write the formula oi potassium carbofiate. This salt
contains the potassium radical K* combined with the carbonate
radical CO 3", and to get equivalence of the dots and dashes
we must take two potassium radicals for one carbonate radical.
We thus obtain the formula K'gCOg" for potassium carbonate.

If it is desired to write the formula of calcium phosphate^ we

62 INORGANIC CHEMISTRY

first of all take the calcium radical Ca" with its two dots, and
the phosphate radical PO4'" with its three dashes. To get the
number of dots and dashes to correspond, we have to take
three calcium radicals and two phosphate radicals. This
gives us 3 X 2 positive charges for the three calcium radicals,
and 2x3 negative charges for the two phosphate radicals, the six
positive charges thus balancing the six negative charges. The
formula then is Ca"3(P04'")2.

It will be observed that in the table we have grouped the
radicals according to the number of dots or dashes associated
with them. Radicals with one dot or dash are called univalent
or monad radicals ; radicals with two dots or dashes are called
bivalent or dyad radicals ; radicals with three dots or dashes
are called trivalent or triad radicals. The monad radicals are
all equivalent to one another, the dyad radicals are also
equivalent amongst themselves, and so likewise are the triad
radicals. In each class, therefore, we have equivalence ; and if
the positive and the negative radicals in a salt belong to the
same class, they combine in equal numbers. If the positive
and negative radicals belong to different classes, then we have
the following rules, of which examples have just been given.
Two monad radicals combine with one dyad radical; three
monad radicals combine with one triad radical ; and three dyad
radicals combine with two triad radicals. The beginner is
strongly advised to write the formula for all possible com-
binations of positive and negative radicals, so as to gain
expertness in writing correct formulae for acids, bases, and
salts. He should at first make use of the table, but in doing
so should endeavour to get the names, symbols, and the
electric charges expressed by the dots and dashes firmly fixed
in his memory, so that he may be able to write, without
consulting the table, the correct formula for any salt whose
systematic name is given to him.

It will be seen that in the table the hydrogen radical
has been separated from the other positive radicals, and that
the hydroxide radical has been separated from the other
negative radicals. This is on account of the peculiar character
of compounds containing hydrogen as the positive radical on
the one hand, and hydroxide as the negative radical on the
other. Compounds which contain only hydrogen as the
positive radical are acids; compounds which contain only

POSITIVE AND NEGATIVE RADICALS 63

hydroxide as the negative radical are basses. All other com-
binations of positive and negative radicals are salts. When
the two peculiar radicals, hydrogen and hydroxide, are com-
bined together we get a substance which, although common,
has chemically speaking very peculiar properties, namely,
water.

According to our system, a substance having the formula
HCl should be called hydrogen chloride, and a substance
having the formula H2SO4, hydrogen sulphate. These names
are sometimes employed, and similarly other compounds
having hydrogen as the positive radical are referred to as
hydrogen salts ; but as a rule an older system is adhered to,
and hydrogen salts according to it are called acids. Thus,
instead of hydrogen chloride we say hydrochloric acid, instead
of hydrogen sulphate we say sulphuric acid. There is a
definite mode of connecting the names of the various acids
with the names of the salts derived from them. If the name
of the negative radical ends in ite the name of the acid
ends in ons. If the name of the negative radical ends in
ate the name of the acid ends in ic. Finally, if the name
of the salt radical ends in ide the name of the corresponding
acid begins with hydro and ends with ic. We have thus the
following correspondence : —

Acid Salt

ous ite

ic ate

Hydro — -ic ide

Corresponding to the sulpha/^j we have sulphur/V acid;
corresponding to the sulph/V^j we have sulphur^wj* acid;
corresponding to the sulph/^/i?j we have hydro^yjX'^Mxic acid.
QWoxides are derived from hydro<^oxic acid, \yiomides from
hydrohxomic acid, \odddes from hydriodSc acid. Exceptions to
this rule of correspondence sometimes occur, but they are not
very frequent.

Occasionally we find that one and the same metal occurs in the
form of two different metallic radicals. Thus mercury is the sole
constituent of the mercurous radical Hg* and of the mercuric
radical Hg". The salts derived from these different radicals have
entirely different properties, and, as we see, are called by differ-
ent names. In such cases one of the names ends in ous, the

64 INORGANIC CHEMISTRY

other ends in ic. In giving these terminations to the different
radicals the following rule is always observed : — The radical
with the smallest number of dots receives the termination ous,
the radical with the greatest number of dots receives the
termination ic. This is exemplified in the case of mercury,
and we see the same thing in the case of iron. The ferrous
radical Fe" has two dots associated with it, the ferric radical
Fe*** has three dots.

The positive radicals may be subdivided into the radical
hydrogen and the metallic radicals which constitute all the
other positive radicals. Similarly, the negative radicals may
be divided into the hydroxide radical and all the other
negative radicals, which are classed together as salt radicals.
Thus we may define acids, bases, salts, and water according
to their radicals as follows : —

Hydrogen radical. Hydroxide radical . Water

Hydrogen radical, Salt radical . Acid

Metallic radical. Hydroxide radical . Base

Metallic radical, Salt radical Salt

We sometimes find that the positive portion of a salt does
not consist of one radical only, but of various radicals com-
bined with the same negative portion. Thus we can have the salt
sodium potassium sulphate Na*K*S04", which contains as the
positive portion both sodium and potassium radicals. Such
a salt is called a mixed salt, and the systematic name is given
by mentioning the two positive radicals first and then the
negative radical. Another instance is found in magnesium
ammonium phosphate Mg"NH4*P04'". In such mixed salts
the rule about dots and dashes holds good, as may be seen by
reference to the formulae. Salts of the opposite type, with one
metallic radical and two or more salt radicals, are not common.

When the positive portion of a compound consists partly
of the hydrogen radical, the compound is said to be an acid
salt. As an example we may take the common bisulphate
of soda^ which has the formula Na'H'SO^", and is named
according to our systematic method sodium hydrogen sulphate.

The phosphates show many instances of mixed salts and
acid salts. We have already seen the mixed salt magnesium
ammonium phosphate Mg"NH4'P04'". The ordinary sodium
phosphate is, in reality, an acid salt ; its formula is Na*2H-P04'",

POSITIVE AND NEGATIVE RADICALS 65

and its systematic name disodium hydrogen phosphate. Anoi her
salt exists which has the formula Na'H'gPO^'", which can
be distinguished from the previous one by naming it sodium
dihyrogen phosphate. We may even have a salt which is
at once the mixed salt and an acid salt — thus, sodium
ammonium hydrogen phosphate Na*N H 4 • H 'PO 4"'.

When a compound consists of one metallic radical com-
bined with two or more negative radicals, one of which is the
hydroxide radical, it is said to be a basic salt. Thus we have
the basic mercuric nitrate Hg"(0H')(N03'), one of the
negative radicals in which is the hydroxide group, and the
other the nitrate group.

Salts of the ordinary type, which consist of metallic radicals
only combined with salt radicals only, are said to be normal
salts.

From a consideration of the formulae, it is evident that
acid salts are intermediate in composition between the acid
and the normal salt of the acid. Sodium hydrogen sulphate
Na'H'SO^" is evidently intermediate between normal sodium
sulphate Na'gSOi" and sulphuric acid H'2S04". These
acid salts are only found when the acid has two or more
hydrogen radicals in its formula. When this is the case, the
hydrogen may be replaced by a metallic radical in succes-
sive steps. For instance, phosphoric acid, with the formula
H'3P04'", has three hydrogen radicals, which can be replaced
in three successive stages by metallic radicals, such as the
sodium radical, with production of the two acid salts
Na-H-2P04'", Na-gH-POi'", and the normal salt Na-3P04'".

Acids which have only one hydrogen radical in the formula
are said to be monobasic ; acids with two hydrogen radicals
are dibasic; acids with three hydrogen radicals are tribasic.
Thus hydrochloric and nitric acids are monobasic ; sulphuric
acid is dibasic ; and phosphoric acid tribasic.

It must not be supposed that salts which are in this formal
sense acid salts, have always the properties of acids. It is true
that sodium hydrogen sulphate is a strongly acid substance,
being sour to the taste, and its solution at once turning blue
litmus red, but this is not invariably the case. The ordinary
sodium phosphate Na*2H'P04'", although formally an acid
salt, is not acid to litmus ; in fact, it has a slightly alkaline
reaction. We must then carefully distinguish between the

66 INORGANIC CHEMISTRY

classification of salts in a purely formal way into acid, normal,
and basic salts ; and the classification of the same substances
according to their action on an indicator like litmus. A
formally acid salt may be acid, neutral, or even alkaline to
litmus ; and, similarly, a normal salt may be acid, alkaline, or
neutral to litmus. When we are dealing with strong bases,
such as sodium hydroxide, potassium hydroxide, or calcium
hydroxide, and with strong acids such as hydrochloric, nitric,
or sulphuric acids, the two modes of classification coincide :
normal salts are then neutral to litmus, and acid salts are
acid to litmus. But when salts formed from strong bases
and weak acids, or weak bases and strong acids, are in ques-
tion, then the two classifications do not coincide. Normal
salts formed from a strong acid and a weak base frequently
give solutions which are acid to litmus. Normal salts which
are formed from a strong base and a weak acid frequently
give solutions which are alkaline to litmus.

CHAPTER XII

DOUBLE DECOMPOSITION

A VERY frequent type of chemical action occurs in aqueous
solution by positive and negative radicals changing places.
In this way what we call a doable decomposition is brought
about, and already in the phenomenon of neutralisation we
have had examples of this kind of action. When we write
the equations for th6 neutralisation of some of the simple
acids by some of the simple bases, indicating the positive
and negative radicals, we find that the neutralisation consists
in the hydrogen radical of the acid combining with the
hydroxyl radical of the base to form water.

Na-OH'

H-cr = Na-cr

HOH'

2K-OH'

H-gSO/ = K-gSO/

2HOH'

Ca"(0H')2

2H-N08' = Ca-(N03')4

2HOH'

A1-(0H')3

3H-Cr = AI-CI3'

3HOH'

SNH^-OH'

H'sPO/" = (NH.OsPO/

" +

3HOH'

The hydrogen radical of the acid changes places with the
metallic radical of the base, and the salt radical of the acid
changes places with the hydroxide radical or hydroxyl of the
base. The positive and negative radicals, therefore, merely
change partners, and in every instance of neutralisation it will
be seen that water is an essential product.

Now, by writing down the formulae of any two salts, and
interchanging their positive and negative radicals, we obtain
an equation which is chemically a possible equation, but
which may or may not express an actual chemical reaction.
There are some rules, however, which we can use for our
guidance in determining which of these double decompositions
actually take place and which do not. It can be said that,
without exception, if one of the positive radicals is hydrogen,
and if one of the negative radicals is hydroxyl, a double

68 INORGANIC CHEMISTRY

decomposition will undoubtedly take place, the hydrogen and
the hydroxyl combining to form water. Similarly, there are
other pairs of positive and negative radicals which, when
brought together, always combine with each other, and thus
bring about double decomposition between salts. In general,
when by the interchange of the radicals of two salts there
is possibility of formation of an insoluble salt, then a double
decomposition will take place. It is therefore of the utmost
importance that the student should know which salts are
soluble in water, and which salts are insoluble in water ; for
on this knowledge depends his power of predicting whether
a possible double decomposition between salts will take place
or not. The following general statements, then, regarding
solubility will be found to be of use : —

All normal nitrates, acetates, diXid' chlorates are soluble

in water.
All potassium, sodium, and ammonium salts are soluble in

water.
All normal chlorides are soluble, except HgCl, AgCl, and

PbCljj.
All normal sulphates are soluble except BaS04, SrS04,

PbS04, and CaSO^.
The ordinary acids are soluble.
All normal phosphates and carbonates are insoluble except

those of potassium, sodium, and ammonium.
All hydroxides are insoluble except those of sodium,

potassium, and ammonium, which are freely soluble,

and those of calcium, strontium, and barium, which

are sparingly soluble.

If we consider the equation —

Na-Cr + K-NOg' = K-Cr + Na'NOs'

Sodium chloride Potassium nitrate Potassium chloride Sodium nitrate

we find that all the substances concerned in it are soluble,
and the above rule about insoluble substances does not there-
fore apply. // must not be supposed, however, that because the
rule does not apply the reaction cannot take place. When the rule
is not applicable, we are simply not in a position to tell whether
the reaction can occur or not.

DOUBLE DECOMPOSITION 69

If we now consider the equation —

Zn-SO^" + Ba"(N03')a = Zn"(N03')2 + Ba'S04"
Zinc sulphate Barium nitrate Zinc nitrate Barium sulphate

we can say that the action, according to the above rule, will
occur, because one of the products of the reaction is insoluble —
that is, the substances on the left of the equation by interchange
of radicals can produce the insoluble barium sulphate. The
action will therefore take place.

Again, if we ask : Will zinc nitrate and sodium phosphate
enter into double decomposition ? we can say that they will,
for zinc nitrate and sodium phosphate are soluble salts, but
by interchange of radicals can produce sodium nitrate which
is soluble and zinc phosphate which is insoluble.

3Zn- (NOaOa + 2Na3-P04'" = Zn-3(POn2 + 6Na-N03'

Zinc phosphate

Suppose now we ask: Will barium sulphate and sodium
chloride enter into double decomposition to produce barium
chloride and sodium sulphate? It is evident that they will
not, for here we should have an insoluble and a soluble sub-
stance reacting together to produce two soluble substances,
a reaction which in view of the above rule is impossible,
for these soluble substances would themselves interact to
produce sodium chloride and the insoluble barium sulphate.

We can therefore reverse the rule, and say that an insoluble
and a soluble salt will not react to produce two soluble salts.

Take now the following case : —

Na-aCOa" + Ba-SO/ = Na-2S04" + Ba'COa"

On one side we have the insoluble salt barium sulphate, on-
the other side we have the insoluble salt barium carbonate.
The conditions for the application of the rule are therefore
not fulfilled, and we cannot say whether the action will take
place or not.

It must be understood that the above rule applies only to
salts^ and not to reactions involving acids. For example, if
we consider the reaction —

Zn"3(PO4"0 2 + 6HC1 = sZn-Cla" + 2H-3P04'"

we should say, if the rule were applicable to this case, that

70 INORGANIC CHEMISTRY

the zinc phosphate would not react with the hydrochloric acid
to produce zinc chloride and phosphoric acid ; but as a matter
of fact, these substances do interact, and the zinc phosphate
dissolves up entirely when treated with hydrochloric acid.
When an acid, therefore, takes part in a double decomposition,
we must be prepared for exceptions to the rule given above.

The displacement of one acid from its salts by another acid,
of which the above reaction is an example, is of great import-
ance. If we treat a solution of sodium chloride with nitric
acid, some of the hydrochloric acid is displaced by the nitric
acid, according to the equation —

Na-Cr + HNO3' = NaNOg' + HCl'

Here we have the same kind of interchange of radicals as
before ; the hydrogen radical, which was originally connected
with the chloride radical, being now combined with the nitrate
radical. This reaction is a balanced action, for if we add
hydrochloric acid to a solution of a nitrate, the following
interchange occurs: —

HCr + Na-NOa' = Na'Cl' + H'NOs'

Here the nitric acid is displaced from its salt by the hydro-
chloric acid.

This kind of reaction between one acid and the salt of
another acid is quite general in aqueous solutions, and in
every case the action is a balanced one when all the sub-
stances are soluble. Any acid is capable of displacing to a
greater or less extent any other acid from its salts, and we
can conveniently define the inherent strength of an acid by
the extent to which it will turn out other acids from solutions
of their salts. It must be clearly understood, however, that
the two acids must compete against each other under con-
ditions equally favourable to both. If one of the acids is
removed from the sphere of reaction, then the other acid is
of course favoured, for the first acid after its removal from the
sphere of reaction is no longer in a position to compete with
the other acid. Now, substances may be removed from the
sphere of reaction in two ways. First, they may separate out
as insoluble precipitates ; and, second, they may separate out
as gases. As most of the common acids are freely soluble

DOUBLE DECOMPOSITION 71

in water, the second mode of removal of the acids is
of greater practical importance than the first in determin-
ing the conditions for the displacement of one acid by
another.

If we heat a non-volatile acid together with salt of another
acid, which is itself volatile, then the second will pass off as
vapour as soon as it is formed, and the result will be that
the non-volatile acid, if taken in sufficient quantity, will
eventually displace all the volatile acid from its salt. Thus,
when concentrated sulphuric acid is warmed with sodium
chloride, it displaces some of the hydrochloric acid from the
chloride, and the hydrochloric acid being in the form of a gas
at the temperature of the experiment, is removed from the
sphere of action as soon as formed ; so that if the action of
the sulphuric acid on the sodium chloride is sufficiently pro-
longed, all the hydrochloric acid will be expelled, according to
the equations —

NaCl + H2SO4 = NaHS04 + HCl

Sodium hydrogen
sulphate

NaCl + NaHSO^ == NagSO* +HC1

Normal sodium
sulphate

The expulsion here takes place in two stages. The tempera-
ture necessary for the first stage is comparatively low, the
temperature necessary for the second stage is much higher;
but in each case the hydrochloric acid is much more volatile
than the sulphuric acid, and is consequently expelled as gas,
this fact being conveniently indicated by the arrow attached
to the formula.

In the above instance the sulphuric acid and the hydro-
chloric acid are not competing on equal terms, because the
hydrochloric acid cannot remain in the sphere of reaction.
Suppose, now, we add sulphuric acid to a dilute solution of
sodium chloride at the ordinary temperature. In that case,
neither the sulphuric acid nor the hydrochloric acid is volatile,
and the hydrochloric acid produced remains in the solution
and consequently continues to compete against the sulphuric
acid* When we consider the extent to \^hicb the sulphuric

72 INORGANIC CHEMISTRY

acid displaces the hydrochloric acid under such conditions,
we find that the sulphuric acid, instead of being inherently
stronger than the hydrochloric acid, as is generally supposed,
is in reality considerably weaker.

The expulsion of nitric acid from a nitrate by means of
sulphuric acid is similar. At the temperature at which the
reaction is conducted, the nitric acid is volatile, and the sul-
phuric acid is not. The consequence is, that if sufficient
sulphuric acid is employed, all the nitric acid is expelled,
according to the equation —

NaNOa' + H-2S04" = NaHSO/ + HNO3'

Sodium hydrogen
sulphate

When the two acids are made to compete against each other
in aqueous solution, under conditions equally favourable to
both, it is found that the nitric acid is considerably stronger
than the sulphuric acid.

When the strength of the acids is measured in this way, by
finding to what extent they displace each other in aqueous
solution, it appears that hydrochloric and nitric acids are about
equally strong, that sulphuric acid is somewhat weaker than
these, and that phosphoric acid is weaker still. Acetic acid
(the acid in vinegar) is so weak in comparison with these
others, that nearly all of it is displaced from its salts by
them. Thus :

NaCgHaOa' + HO = Na-Cl' + H-CgHgOg'
Sodium acetate Hydrochloric acid Sodium chloride Acetic acid

Acetic acid may be taken as a typical weak acid, the acid
properties of which still remain quite pronounced. For
example, acetic acid, although weak, has still a very sour taste,
and at once reddens litmus.

Carbonic add H'gCOs" is a very much weaker acid than
acetic acidl Its solutions have no sour taste, and are just
capable of reddening litmus. Sulphuretted hydrogen^ or hydro-
sulphuric acid, H'gS", is another acid of much the same
strength as carbonic acid. These two acids are also similar
in another respect — namely, they are not very soluble in water,
^nd are easily expelled from aqueous solution by heating, the

DOUBLE DECOMPOSITION 73

carbonic acid breaking up at the same time into carbon
dioxide and water, according to the equation —

H'aCOs' = HsO + CO2

When solutions of the carbonates are treated with hydrochloric,
or even with acetic acid, the carbonates are converted into
chlorides or acetates almost completely, according to the
equations —

NagCOs + 2HCI = 2NaCl + H^O + COg

K2CO3 + 2HC2H3O2' = 2 KCgHsOa' + H2O + CO2

These reactions are favoured not only by the inherent weak-
ness of the carbonic acid, but also by its tendency to split
up into water and carbon dioxide, which passes off as gas.
Similarly, if a soluble sulphide is treated with hydrochloric,
nitric, sulphuric, or acetic acid, sulphuretted hydrogen is at
once produced, on account both of its inherent weakness, and
on account of its being a gas which is incapable of remaining
to any great extent in solution.

NagS + aH-CgHgOg' = 2Na-C2H302' + HgS

When solid insoluble carbonates are treated with any strong
acid, they split up with evolution of carbon dioxide. For
example, if calcium carbonate is treated with hydrochloric acid,
though the calcium carbonate is in the dense form of marble,
it is at once decomposed, according to the equation —

CaCOs + 2HCI = CaClg + U^O + CO2

Even acetic acid is a sufficiently strong acid to decompose
solid carbonates in this way.

Some insoluble sulphides can be similarly decomposed by
strong acids. For example, ferrous sulphide, FeS, although
insoluble in water, is easily attacked by moderately dilute
hydrochloric acid, with evolution of sulphuretted hydrogen,
according to the equation —

FeS + 2HCI = FeCl2 + H2S

Ferrous sulphide Ferrous chloride Sulphuretted hydrdgen

This action is usually made use of in the laboratory to prepare

74 INORGANIC CHEMISTRY

sulphuretted hydrogen gas. Some insoluble sulphides, how-
ever, resist decomposition by acid. For example, mercuric
sulphide HgS, and arsenious sulphide AsgSg, are scarcely
attacked by a solution of hydrochloric acid, and very little
sulphuretted hydrogen is expelled from them.

The displacement of bases by each other from their salts
is subject to the same rules as those we have just seen regard-
ing acids. If two bases compete with each other under fair
conditions, their strength can be measured by the extent to
which they can displace each other. Thus it is found that
potassium hydroxide and sodium hydroxide are almost equally
strong. If one of the bases is removed as soon as it is formed,
then, of course, the competition does not take place under
fair conditions, and so one base may throw out another base
from its salts, although the second is as strong as the first.
Thus, if we add a concentrated solution of potassium hydrox-
ide to a concentrated solution of calcium chloride, calcium
hydroxide will be produced according to the equation —

2KOH + CaCl2 = 2KCI + Ca(OH)2

Potassium hydroxide Calcium chloride Potassium chloride Calcium

hydroxide

Now, this calcium hydroxide is only sparingly soluble in water,
and consequently soon begins to separate out. The portion
which is separated out as solid has practically left the sphere
of the reaction, and so it is possible by means of potassium
hydroxide to displace a very large proportion of calcium
hydroxide from its salts, if the solutions are concentrated. If
enough water is present, however, to dissolve all the calcium
hydroxide which would be produced — that is, if we work with
dilute solutions, then it is found that calcium hydroxide is
almost as strong a base as the potassium hydroxide.

As most of the hydroxides are insoluble or sparingly soluble
in water, sodium or potassium hydroxide, which are soluble,
can in general displace them from their salts. The following
are examples of such reactions : —

ZnS04 + 2NaOH = NaaSQ^ + Zn(0H)2

Zinc hydroxide

CuCL + 2KOH = 2KCI + Cu(0H)2

Cupric hydroxide

Ammonium hydroxide is a base which bears much the same

DOUBLE DECOMPOSITION 75

relation, so far as strength is concerned, to potassium and
sodium hydroxide, as acetic acid does to nitric and hydrochloric
acid. Consequently, if we add caustic soda to a solution of
an ammonium salt, even though all the ammonium hydroxide
remains in solution, a very large proportion of it will be
displaced from its salt. If the solution is boiled, the
ammonium hydroxide splits up into water and ammonia gas,
which leaves the solution, so that, both on account of its
inherent weakness, and of its volatility, ammonium hydroxide
can always be displaced from ammonium salts by heating the
solutions with caustic alkalies.

(NH4)2S04 + 2NaOH = NagSO^ + 2NH4OH

NH4OH = NH3 + H3O

Ammonia

CHAPTER XIII

lONISATION AND DISPLACEMENT OF RADICALS

When an acid, base, or salt is dissolved in water, it is, accord-
ing to modern theory, split up to some extent into its positive
radical and its negative radical. The extent to which this
splitting up occurs varies very greatly in different cases, but
it may be noted at once that all ordinary salts, all strong
bases, and all strong acids are in great measure split up by
water into their positive and negative radicals, and that the
more as the solution becomes more dilute. Weak acids and
weak bases are much less decomposed by water into their
positive and negative components than the salts formed from
them.

The free positive and negative radicals thus produced by
the action of water on the salts, acids, and bases are called
ions, and a consideration of the electrical properties of the
aqueous solutions shows that the ions are charged with
positive and negative electricity in accordance with the dots
and dashes attached to the radicals in the symbols given on
page 59. If we take the gram as our unit of weight, so
that the symbols express a given number of grams, then each
dot stands for 96,500 coulombs of positive electricity, and
each dash for an equal amount of negative electricity. Thus
we can say that i gram of hydrogen existing in aqueous
solution as an ion has a charge of 96,500 coulombs of
positive electricity since the symbol H* expresses i gram, and
the dot attached to it expresses this amount of electricity.
Similarly, the symbol SO 4" expresses that 96 grams of sulphate
radical existing as an ion in aqueous solution is charged with
twice 96,500 coulombs of negative electricity.

In an ordinary solution of common salt we may take it that
somewhat more than one-half of the total sodium chloride

lONISATION AND DISPLACEMENT OF RADICALS 77

present in the solution is split up or ionised into the sodium
ion Na* and the chloride ion Cr. The degree of ionisation is
greater or less than this according to the concentration of
the solution. It is invariably the case that as concentration
increases ionisation diminishes. If we wish, therefore, a dis-
solved salt to be almost entirely split up into its constituent
ions, we take a dilute solution of that substance. In a very
dilute solution of sodium chloride — a solution containing, say,
one-tenth of a gram of sodium chloride dissolved in a litre
of water — the ionisation is practically complete.

The simple radicals can all exist in the free state as elements.
The compound radicals are, as a rule, incapable of independent
existence except as charged ions. We find that the tendency
of the elements to assume the form of electrically charged
radicals varies very much, some elements having a great ten-
dency to become charged with electricity and pass into solution
as ions, whilst other elements have little or no tendency in
that direction. If we take a piece of metallic iron and im-
merse it in a solution of copper sulphate, the iron passes into
solution and metallic copper is precipitated, the chemical
action being expressed by the equation —

Fe + CU-SO4" = Fe-S04" + Cu

Iron Copper sulphate Ferrous sulphate Copper

The iron has here displaced the copper from the solution
of copper salt. In other words, the iron has assumed the
form of a charged radical, whilst the copper has lost its electric
charge and become metallic copper. We can, therefore, infer
that iron has a greater tendency to assume an electric charge
in solution than copper has. Similarly, if we immerse a strip of
copper in a solution of a mercuric salt, say mercuric chloride,
the copper, immediately becomes coated with a grey film of
metallic mercury, and the originally colourless solution assumes
a blue colour, showing that a copper salt is now in solution.
The action which has actually taken place is as follows : —

Cu + Hg-Cl'a = Cu-Cl'a + Hg
Copper Mercuric chloride Copper chloride Mercury

The mercuric radical has lost its charge and become metallic

78 INORGANIC CHEMISTRY

mercury, whilst the metallic copper has taken up the charge
originally on the mercury, and become the copper radical.

Considering the various metals from the point of view that
they have each a perfectly definite tendency to assume elec-
trical charges, we find that if we bring them into intimate contact
with solutions of one another's salts, they displace each other
in a perfectly definite order, the metal with the greater tendency
to unite with electricity passing into solution as a charged ion,
• and the metal with the smaller tendency to unite with elec-
tricity being thrown out of solution as the uncharged metal.
The order in which some of the commoner metals replace each
other from solutions of their salts in this way is given in the
following table : —

Sodium

Magnesium

Aluminium

Zinc

Iron

Tin

(Hydrogen)

Copper

Mercury

Silver

Gold

The metal sodium, which heads the list, has the greatest ten-
dency to unite with positive electricity, and displaces all the
other metals from their salt solutions. It should be noted that
by salt solutions we generally understand either chlorides or
sulphates, in order that complications which sometimes arise
in connection with other negative radicals may be avoided.
In general, we may say that any metal which occurs on this list
will displace from their salt solutions the metals which follow
it in the list, but will not displace the metals which precede it,
being rather displaced from its own solutions by these metals.
We can tell at once, therefore, from an inspection of
this table, that metallic iron will turn out copper from
copper sulphate, but that it will not turn out zinc from zinc
sulphate.

Hydrogen, which has a peculiar position in the list of positive
radicals, has also a peculiar position in this electro-chemical

lONISATION AND DISPLACEMENT OF RADICALS 79

list. On account of its being a gas it cannot very readily be
brought into contact with aqueous solutions, and so cannot
readily be made to displace the metals which occur below it
in the list. It can, however, be readily enough displaced
from hydrogen salts — that is, the acids such as hydrogen
chloride (hydrochloric acid), or hydrogen sulphate (sulphuric
acid) — by means of metals which precede it in the list, and
cannot under any conditions be displaced from these acids by
the metals which follow it in the list.

We can thus see at once from the table that a metal such
as zinc or iron will be acted upon by hydrochloric or dilute
sulphuric acid with evolution of hydrogen, and that a metal
like copper or mercury will not be attacked by these acids.
In the first case, the iron and zinc have a greater tendency to
assume electrical charges and become positive radicals than
the hydrogen has, and, in the second case, the hydrogen has a
greater tendency to remain combined with positive electricity
than either the copper or the mercury, and consequently will
not be displaced from its solutions by these metals.

In connection with this we may consider the action of the
various metals on water. Any metal which precedes hydrogen
in the list is capable of attacking water at some temperature ;
sodium, for example, at once liberates hydrogen from water
according to the equation —

2Na + 2HOH = 2NaOH + Hg

Sodium Water Sodium hydroxide Hydrogen

The Other metals preceding hydrogen in the list do not attack
water at the ordinary temperature, but if they are heated with
water to a sufficiently high degree, the water is attacked with
formation of a metallic oxide or hydroxide and liberation of
hydrogen — e.g,

2AI + 3H2O = AI2O3 + 3H2

Aluminium Aluminium oxide Hydrogen

2AI + 6H2O = 2A1(0H)3 + 3H2

Aluminium hydroxide

None of the metals which follow hydrogen in the list are capable
of attacking water at any temperature.

8o INORGANIC CHEMISTRY

When we consider that any substance found in the earth's
crust is now, or has been at some period, in contact with water
at a high temperature, we see that we cannot expect the metals
which precede hydrogen in the list to occur in the free state in
nature, whilst we may very reasonably expect that the metals
which follow hydrogen in the list should be found in nature in
the metallic state. This is actually the case : sodium, magnes-
ium, aluminium, and zinc are never found as metals, but always
combined with oxygen, or sometimes with sulphur. Iron only
occurs in nature in minute quantities in some rocks, or «lse in
meteorites, which, properly speaking, are not a part of the
earth's crust ^t all.

On the other hand, copper, mercury, silver, and gold, are
found in considerable quantities in the metallic state, although
(with the exception of gold) they are also found combined
with oxygen and with sulphur.

The chief elements which can exist as negative ions in
solution are chlorine, bromine, and iodine, which form the
negative radicals of chlorides, bromides, and iodides. Just
as we ^ can arrange the positive elements, so we can arrange
these three negative elements in the order in which they can
displace each other. The list is as follows : —

Chlorine
Bromine
Iodine

Chlorine is capable of displacing bromine from bromides and
iodine from iodides ; bromine is capable of displacing iodine
from iodides, but not chlorine from chlorides. Iodine is
incapable of displacing either bromine from bromides or
chlorine from chlorides. The equations representing the
various actions which can occur in the case of the sodium
salts are as follows : —

C12

2Na-Br'

2Na-Cr

Br2

Bra

2Nar

2NaBr'

2Nar

2Na-Cr

Chlorine, bromine, and iodine are all capable of displacing
sulphur from soluble sulphides or hydrosulphides, the re-

lONISATION AND DISPLACEMENT OF RADICALS 8i

actions being represented as follows in the case of hydrogen
sulphide : —

Chlorine

HgS

Sulphuretted
hydrogen

2HCI +

Hydrochloric
acid

S
Sulphur

Brg

Bromine

HjS

2HBr +
Hydrobromic acid

Iodine

HjS

2HI +

Hydriodic acid

Chlorine is also capable of replacing oxygen to some extent
from solutions of oxides or hydroxides, although in most
cases the displacement is complicated by the occurrence of
other actions (Chapter XXV. ). These actions are almost
absent when the hydroxide considered is water itself. A solu-
tion of chlorine in water, especially when exposed to sunlight,
decomposes as follows : —

2CI2 +

Chlorine

2H2O

Water

4HCI +

Hydrochloric acid

Oxygen

hydrochloric remaining in the solution, and the oxygen being
liberated as gas. Bromine and iodine are incapable of dis-
placing oxygen from water, although they act readily in a
somewhat complicated way on solutions of soluble hydroxides
(Chapter XXVL).

CHAPTER XIV
ELECTBOLTSIS

It has been stated in the preceding chapter that the sodium
chloride in an ordinary solution of salt is split up to at least
half its extent into electrically charged ions. If the solution
is very dilute, the decomposition into ions is almost com-
plete. We have now to consider the electrical properties of
such a solution. Since the solution, as a whole, is electrically
neutral — i.e, is neither positively nor negatively charged — the
amount of positive electricity carried by the sodium ion
must be exactly equal to the amount of negative electricity
carried by the chloride ion. Suppose now we introduce into
the solution two conductors, one charged with, positive and
the other with negative electricity. Since opposite electrici-
ties attract, the positive radical will be attracted to the
negatively charged conductor, and the negative radical will
be attracted to the positively charged conductor.

Now, we can charge conductors by connecting them with
the opposite terminals of an electric battery. At one ter-
minal there is always a charge of positive electricity ; at the
other terminal there is always a charge of negative electricity,
and electrical conductors attached to these terminals assume
similar charges which are maintained by the action of the
battery. If, then, we immerse in a solution of common salt
conductors connected with the terminals of a galvanic
battery, the positive ion will travel towards the negative
terminal, and the negative ion will travel towards the positive
terminal. There will thus, on the whole, be separation of the
positive radical of the salt from the negative radical.

We may now ask what is likely to happen when the charged
negative radical actually comes in contact with the positively
charged conductor. The opposite charges neutralise each
other ; the negative radical thus loses its charge and becomes
electrically neutral. Now, when any radical loses its charge
of electricity its chemical properties are profoundly modified,

ELECTROLYSIS 83

So long as the chloride radical in an aqueous solution of
sodium chloride retains its negative charge and exists as
chloride ion, it has certain well-marked properties, use of
which is made in the general tests employed for detecting
soluble chlorides. As soon as it loses its charge and ceases
to be a negative ion it assumes altogether different properties :
two of the uncharged radicals unite to form the free element
chlorine, which we represent by the formula Clg.

2Cr - CI2
Chloride ion Chlorine

This chlorine is a greenish-yellow gas, only moderately soluble
in water, and possessed of a pungent, suffocating smell, by means
of which it is easily recognised. It differs altogether, there-
fore, from the chloride ion, which is colourless and odourless.

When the sodium radical which exists in the solution of
sodium \chloride as a positive ion loses its charge on coming
in contact with the negatively charged electrode, it becomes
ordinary metallic sodium. Now, as we have seen in the pre-
ceding chapter, metallic sodium cannot exist in contact with
water. It immediately reacts with water, according to the
equation —

2Na + 2H2O = 2NaOH + Hg

caustic soda being produced in the solution, and hydrogen
liberated as gas.

When, therefore, we decompose a solution of sodium
chloride by means oT a current of electricity, chlorine gas
is produced at one of the terminals and hydrogen gas at the
other, the liberation of hydrogen being accompanied by the
formation of caustic soda.

The process of separating the positive and negative ions
from each other, and discharging them by means of an electric
current, is called electrolysis. The conductors connected
with the two terminals of the electric battery are called the
electrodes, that connected with the zinc end of an ordinary
cell being called the negative electrode or kathode, that
connected with the other terminal of the battery being termed
the positive electrode or anode. The negative ion, which
moves towards the positive electrode or anode, is called the
aniQii. The positive ion, which moves towards the negative

INORGANIC CHEMISTRY

electrode or kathode, is called the kation. The solution
which is decomposed is called the electrolyte, but this term
is sometimes applied also to the substance dissolved in the
water of the electrolytic solution.

It is possible, by passing an electric current through an
aqueous solution of an electrolyte, to separate the anion from
the kation, and to make them lose their charges. Sometimes
the discharged anion or kation may be liberated as the
element — for example, in the case of chlorine being evolved
by the discharge of the chloride ion. Sometimes, however.

Wire conneded with
zinc pole of Battery

Negative electrode
or Kathode"

Wire connected with
platinum pole of Battery

*-Na CI-*
Kation Anion

Positive electrode
or Anode

Fig 1 8. — Diagrammatic Representation of Electrolysis of a Solution of

Sodium Chloride.

as in the case of potassium, the discharged ion is not capable
of existing as such under the conditions in which the elec-
trolysis takes place. The discharged ion may then, as
frequently happens, attack the water of the electrolyte. For
example, if we electrolyse a solution of potassium sulphate
the discharged potassium at once attacks the water, accord-
ing to the equation —

2K + 2H2O = 2KOH + Ha

Potassium hydroxide Hydrogen

with formation of potassium hydroxide and hydrogen. The

ELECTROLYSIS 85

discharged sulphate radical is altogether incapable of inde-
pendent existence, and it also attacks the water, according
to the equation —

2SO4 + 2H2O = 2H2SO4 + O2

Sulphuric acid Oxygen

with formation of sulphuric acid and oxygen.

If we only pay attention to the gases which are evolved
during the electrolysis of potassium sulphate, we find that
they are oxygen and hydrogen in the proportions necessary
to form water. It would therefore appear as if the electric
current merely decomposed the water in which the potassium
sulphate was dissolved. If, however, we attend to all the
chemical changes which have taken place during the action,
we find that potassium hydroxide has been produced at the
kathode and sulphuric acid at the anode, as may be readily
shown by adding a little litmus to the solution. Potassium
sulphate before electrolysis is neutral, and on the passage
of a current the purple litmus will be observed to become
blue at the kathode and red at the anode.

When a solution of an ^cid is electrolysed, the discharged
hydrogen usually comes off at the kathode as hydrogen gas.
Thus, a solution of sulphuric acid gives hydrogen at the
kathode and oxygen at the anode, produced by the action of
the discharged sulphate ions, according to the preceding
equation. Here it will be observed that sulphuric acid is
reproduced at the anode, so that the only chemical change
observed is that hydrogen is evolved at the kathode and
oxygen at the anode, the total action being again apparently
a mere decomposition of water by the electric current. A
closer investigation, however, shows that this is not the only
action which takes place. If the anode* and kathode compart-
ments are separated by a porous material which permits free
passage to the current, but does not permit mechanical mixing
of the solutions in the two compartments, it will be found
that after a current has been passed for some time sulphuric
acid will have been transferred from the kathode to the anode.
Thus the electrolysis is not the mere decomposition of water,
but a reaction in which the sulphuric acid dissolved in the
water plays a part.

Similarly, when a solution of sodium hydroxide is electrolysed

86 INORGANIC CHEMISTRY

the total action is decomposition of water, and transference of
sodium hydroxide from the anode to the kathode, the separate
actions being —

At anode . . * . 4OH = 2H2O + Og

Oxygen

At kathode 2Na + 2H2O = 2NaOH + Hg

Sodium hydroxide Hydrogen

When a solution of hydrochloric acid is electrolysed under
proper conditions, hydrogen appears at the kathode and
chlorine at the anode, so that the total action is apparently
decomposition of the hydrochloric acid. If the hydrochloric
acid solution which is electrolysed is dilute, another action
takes place at the anode. Some of the discharged chloride ion,
instead of coming off as chlorine gas, attacks the water of the
electrolyte, with formation of hydrochloric acid and oxygen, in
accordance with the equation —

4CI + 2H2O = 4HCI + O2

Hydrochloric acid Oxygen

We may therefore have two actions going on round the
anode at one and the same time, the result being production
of both chlorine gas and oxygen gas in varying proportions,
according to the conditions under which the electrolysis is
conducted.

In what has preceded, it has been assumed that the material
of the electrodes is not attacked by the discharged ions.
Most metallic conductors may be used as kathodes, because
a discharged kation very seldom attacks them. It is much
more difficult to get an anode which is not attacked. On a
small scale, the anode generally employed is platinum, which
is not attacked readily by any discharged anion. On the
manufacturing scale, the anode almost invariably employed
is a form of carbon which conducts electricity, being practically
the same as the carbons used in arc-lamps (Chapter XXII.).

If we electrolyse a solution of sulphuric acid between copper
electrodes, hydrogen comes off at the copper kathode as usual.
When the sulphate ion, however, is discharged at the anode it
does not attack the water of the electrolyte, as it would were
a platinum anode employed, but rather attacks the copper of

ELECTROLYSIS 87

the anode, the action being one of simple union, as expressed
by the equation —

SO4 + Cu = CUSO4

Copper sulphate

The copper anode thus dissolves up, and copper sulphate is
formed in the solution.

Suppose now we consider the electrolysis of a solution of
copper sulphate between two copper electrodes. As we have
just seen, the copper anode is attacked by the discharged
sulphate ions with production of copper sulphate. On the
other hand, the discharged kation is capable of existence as
metallic copper and is deposited as such on the kathode.
No gas is evolved during this electrolysis, and the whole action
seems to be a transference of copper from the anode to the
kathode, the copper anode diminishing in weight, and the
copper kathode increasing.

Such processes of electrolysis as the above are carried out
on the large scale in the electrolytic refining of metals like
copper, and in the electro-deposition of metals — that is, the
process usually known as electro-plating. The aim of electro-
plating is to protect a metallic object by covering it with a
thin coating of a valuable metal, such as silver or gold, which
is more capable of resisting ordinary atmospheric influences.
The valuable metal, say silver, is taken as the anode, and the
object formed of metal to be silver-plated, say copper, as the
kathode, the solution employed being some salt of silver.
When a current is passed through the solution, the silver
dissolves up and is deposited on the copper object, the con-
ditions of electrolysis being so chosen that the deposit adheres
closely to the kathode in the form of a fine continuous film.
The copper is then said to be silver-plated.

When the same electric current is passed in succession
through a series of solutions, say of sulphuric acid, copper
sulphate, sodium hydroxide, and hydrochloric acid, it is
found that all the electrodes where hydrogen is evolved
{Le. in the sulphuric acid, sodium hydroxide, and hydro-
chloric acid solutions), the quantity of hydrogen produced
is exactly the same. In the solution of copper sulphate,
where no hydrogen is evolved, the quantity of copper deposited
is exactly equivalent to the quantity of hydrogen appearing

88 INORGANIC CHEMISTRY

at the other kathodes. For every gram of hydrogen liber-
ated by a given current, 31.5 grams of copper are deposited.
A reference to the table of atomic weights and of positive
radicals will show that these proportions are equivalent pro-
portions. This equivalence finds a representation in the dots
affixed to the symbols of the positive radicals. A given
current will neutralise the same quantity of positive electricity
in any part of the circuit. The quantity of positive electricity,
therefore, which is neutralised at the various kathodes is in
every case the same. But, since a given quantity of positive
electricity is always associated with the same quantity of
hydrogen, it is evident, therefore, that at the kathodes equal
quantities of hydrogen will be evolved. Further, since 63
grams of copper, according to our table, are associated with
the same amount of positive electricity as 2 grams of hydrogen,
the current which will neutralise the positive electricity of one
gram of hydrogen, will serve to neutralise the positive electricity
of 31.5 grams of copper.

What holds good here for the substances produced at the
kathode also holds good for the substances produced at the
anode. The amount of oxygen evolved at the anode in the
sulphuric acid, copper sulphate, and sodium hydroxide solu-
tions, is exactly the same in the three cases.

Not only is this so, but the substances produced at the
kathode are exactly equivalent to the substances produced at
the anode. Thus, the hydrogen liberated at the kathode in the
sulphuric acid, or caustic soda, is exactly in the proportions
necessary to combine with the oxygen which is liberated by
the same current at the anode.

In the case of the hydrochloric acid solution, both chlorine
and oxygen may be liberated at the anode. Here the amount
of chlorine which is produced is not equivalent to the hydrogen
which is simultaneously produced at the kathode, nor is the
amount of oxygen equivalent to this amount of hydrogen, but
the oxygen and chlorine together are exactly equivalent to the
hydrogen which is liberated by the same current. That is, if
the chlorine is made to combine with hydrogen to produce
hydrochloric acid, then there will be just enough hydrogen
over to combine with the oxygen for the production of water.

Equivalence of the products at the anode and kathode
holds good for the substances produced in the solution as well

ELECTROLYSIS 89

as for the gases evolved. When sodium sulphate is electro-
lysed, sodium hydroxide is formed at the kathode, and sul-
phuric acid at the anode. Now, these two substances are
formed by the same current in exactly equivalent proportions.
The sodium hydroxide produced at the kathode is in precisely
the right proportion for neutralising the sulphuric acid produced
at the anode.

These facts regarding the chemical equivalence of the various
substances produced at the electrodes on electrolysis were first
experimentally ascertained by Faraday, who also showed that a
given amount of electricity passed through a circuit always
liberated the same amount of a given material, e.g. hydrogen,
from an electrolyte, no matter what the substance dissolved in
the water might have been, provided, of course, that hydrogen
was the only gas produced at the kathode.

These facts are generally known under the name of Faraday's
laws of electrolysis, and it will be seen that our system of sym-
bolisation is in accordance with them. Chemically equivalent
quantities — that is, quantities which can either combine with
each other, or which can combine with the same quantity of
other substances — are, according to our system, charged with
equal amounts of electricity as expressed by the dots and
dashes.

The idea that the positive and negative radicals of salts
are to a great extent independent of each other in aqueous
solution affords a ready explanation of the ease with which
double decomposition in aqueous solution can take place
between these substances. The double decompositions, ex-
amples of which have been given in the preceding chapter,
take place practically instantaneously. On the other hand,
reactions in aqueous solutions which are not merely recom-
binations of positive and negative radicals take place com-
paratively slowly.

CHAPTER XV

EXAMPLES OF CHEMICAL TBANSFOBMATION

By applying the general rules which have been given in the
previous chapters concerning acids, bases, and salts, the
student is in a position to solve many practical questions
regarding the conversion of one compound into another.

For example, if he is asked how he might prepare zinc
chloride from zinc hydroxide, he can at once say that zinc
hydroxide, being a base, can be converted into any salt of
zinc by treatment with the corresponding acid ; so that for the
conversion of zinc hydroxide into zinc chloride, the hydroxide
has merely to be dissolved in the requisite amount of hydro-
chloric acid.

Zn(0H)2 + 2HCI = ZnCla + 2H2O

If, again, he is asked how metallic zinc may be converted into
zinc chloride, he knows, from the list on p. 78, that zinc is
capable of displacing hydrogen from acids, so that if metallic
zinc is treated with hydrochloric acid, zinc chloride will be
formed, and hydrogen evolved according to the equation —

Zn + 2HCI = ZnCL + H

If he is asked to convert zinc chloride into zinc sulphate, he
knows that sulphuric acid, being less volatile than hydrochloric
acid, can drive out the latter from its salts on heating. Thus,
to convert zinc chloride into zinc sulphate it is merely neces-
sary to heat the chloride with sulphuric acid, when reaction
will occur as follows : —

ZnClg + H2SO4 = ZnSO^ + 2HCI

hydrochloric acid being expelled as a gas.

EXAMPLES OF CHEMICAL TRANSFORMATION 91

Suppose, however, he is required to convert zinc sulphate
into zinc chloride. This cannot be done by heating zinc sul-
phate with hydrochloric acid, for the hydrochloric acid, being
more volatile, is expelled on heating, the zinc sulphate remaining
behind unchanged. Some other method must therefore be
adopted, and the student might, for example, make use of his
knowledge of the solubility of the various salts. Zinc sulphate and
zinc chloride, according to the rules given on page 68, are both
soluble in water. Now, by interchange of radicals it is possible
in aqueous solution to convert zinc sulphate into zinc chloride,
provided a salt be selected which will give an insoluble com-
pound with the negative radical of zinc sulphate. This salt
must, of course, contain the chloride radical in order to give
zinc chloride. The problem, therefore, resolves itself into
finding a radical whose chloride is soluble and whose sulphate
is insoluble. If the student refers to the table on p. 68, he
finds that barium fulfils these conditions. By adding, there-
fore, a solution of barium chloride to a solution of zinc
sulphate interchange of radicals takes place, with formation
of insoluble barium sulphate and soluble zinc chloride, the
equation being —

ZnS04 + BaClg = BaS04 + ZnCl

The insoluble barium sulphate may be filtered off, and if
the substances were taken in the proportions indicated by the
equation, the solution will contain nothing but zinc chloride.
This type of reaction is quite general. We can convert any
soluble sulphate into the corresponding soluble chloride by
the addition of barium chloride to the solution of the sulphate.

If the chloride which is required from the soluble sulphate
is itself insoluble, we get it at once by double decomposition
with any soluble chloride. Thus, if we have to convert silver
sulphate, which is soluble, into silver chloride, which is in-
soluble, we have merely to add to the silver sulphate solution
a solution of any soluble chloride, say sodium chloride, and
we at once get silver chloride precipitated in accordance with
the equation —

Ag2S04 + 2NaCl = 2AgCl + Na2S04
This silver chloride may be filtered off from the soluble sodium

92 INORGANIC CHEMISTRY

sulphate produced at the same time, and thus obtained free
from the other salts.

Besides this mode of precipitation by means of a suitable
salt, there is another very important method by means of
which soluble sulphates may be converted into soluble chlor-
ides. This method is the one which is most generally used, not
only in this case, but in most similar cases. The method
is to carry out the reaction in two stages. We know that all
carbonates except the alkaline carbonates are insoluble in
water. We can therefore convert any soluble sulphate into
the insoluble carbonate which corresponds to it by adding
to the sulphate solution a solution of alkaline carbonate.
Thus, we can precipitate zinc carbonate from a solution of zinc
sulphate by the addition of sodium carbonate.

ZnSO^ + NaaCOg = ZnCOg + NagSOA

It should be noted in this connection that basic carbonates
are very frequently precipitated by sodium carbonate instead
of the normal carbonates. The precipitate which is actually
obtained by the addition of sodium carbonate to a solution
of zinc sulphate, is not the normal carbonate, as represented
by the above equation, but a basic carbonate which is, like
the normal carbonate, insoluble.

Now, we know that all carbonates are decomposed by acids,
with evolution of carbon dioxide. Thus, when we treat zinc
carbonate with hydrochloric acid, reaction takes place accord-
ing to the equation —

ZnCOa + 2HCI = ZnCla + HgO + CO2

zinc chloride being produced, and carbon dioxide escaping
from the solution as gas. Consequently, by treating a solution
of zinc sulphate with sodium carbonate, filtering off the basic
zinc carbonate produced by the decomposition, and then dis-
solving the latter in hydrochloric acid, we can convert zinc
sulphate into zinc chloride.

The hydroxides, like the carbonates, are nearly all insoluble
in water. We can therefore precipitate hydroxides by means of
sodium hydroxide, instead of carbonates by means of sodium
carbonate. Thus, if we add the requisite amount of sodium

EXAMPLES OF CHEMICAL TRANSFORMATION 93

hydroxide to a solution of zinc sulphate, reaction takes place
as follows : —

ZnS04 + 2NaOH = Zn(0H)2 + Na2S04

The zinc hydroxide thus produced may be filtered off and
separated from the sodium sulphate, and then dissolved in
hydrochloric acid, again with production of zinc chloride.

Zn(0H)2 + 2HCI = ZnCla + 2H2O

Since hydroxides and carbonates are soluble in acids with
equal readiness, it does not in the least matter whether the
carbonate which is precipitated by sodium carbonate is the
normal or the basic carbonate, which is intermediate between
the hydroxide and the normal salt; for in either case, on
treatment with acid the corresponding salt will be formed,
the only difference being in the amounts of water and of
carbon dioxide produced at the same time.

It sometimes happens that the metallic radical considered
has no insoluble compounds with salt radicals. Thus, if the
problem before the student is to convert sodium nitrate into
sodium chloride, no precipitation method in the above sense
is available, because all the compounds of sodium with salt
radicals are soluble, and all the compounds of the nitrate
radical with metallic radicals are soluble. We cannot, there-
fore, add to a dilute solution of sodium nitrate any salt which
will precipitate either the sodium or the nitrate radical — that
is, we cannot bring about double decomposition by the produc-
tion of an insoluble substance. Even in such cases, however,
a precipitation method may be used if we deal with very strong
solutions and properly selected temperatures. Thus it is pos-
sible, by bringing sodium nitrate and potassium chloride
together in presence of only a little hot water, to get double
decomposition to occur, according to the equation —

NaNOa + KCl = NaCl + KNOg

sodium chloride falling out of the solution. As will be seen
in the chapter on potassium, this method is actually adopted,
not for the conversion of sodium nitrate into sodium chloride,
but for the production of potassium nitrate.
Another method, in this case an indirect one, might be

94 INORGANIC CHEMISTRY

adopted to convert sodium nitrate into sodium chloride. If
the sodium nitrate is subjected to electrolysis, sodium hy-
droxide is formed at the kathode. The solution of sodium
hydroxide thus obtained could be neutralised by means of
hydrochloric acid, and thus sodium chloride would be pro-
duced.

If the problem before the student is to convert metallic
copper into copper chloride CUCI2, he might adopt several
methods, all more or less indirect, since it is impossible to
obtain copper chloride directly from metallic copper by the
action of hydrochloric acid alone. One method would be to
heat the metallic copper in air, so as to obtain copper oxide
CuO. This oxide being ba«c, dissolves in hydrochloric acid,
with formation of copper chloride according to the equation —

CuO + 2HCI = CuClg + HgO

Again, although copper is not soluble in dilute hydrochloric
acid, it is soluble in hot concentrated sulphuric acid, copper
sulphate being produced in accordance with the equation —

Cu + 2H2SO4 = CUSO4 + SO2 + 2H2O

This copper sulphate could then be converted into copper
chloride by any of the methods already given for the conver-
sion of zinc sulphate into zinc chloride. Thus it might be
precipitated by means of sodium carbonate, and the basic
carbonate so obtained dissolved in hydrochloric acid.

Instead of sulphuric acid, nitric acid might be employed as
the solvent for metallic copper, the action then being —

3CU + 8HNO3 = 3Cu(N03)2 + 2NO + 4H2O

The copper nitrate produced in this way could either be con-
verted into chloride in the same way as the sulphate, or it
might be subjected to heat, when it decomposes as follows : —

2Cu(N08)2 = 2CuO + 4NO2 + O2

with formation of copper oxide, which could then be dissolved
in hydrochloric acid, as in the preceding instance.

When the substance to be dealt with is insoluble in acids,
there is sometimes considerable difficulty in converting it into
soluble compounds with the same metallic radical, For example,

EXAMPLES OF CHEMICAL TRANSFORMATION 95

barium sulphate is not only insoluble in water, but also in
the common acids. If it is desired, then, to convert barium
sulphate into barium chloride, we cannot employ any of the
precipitation methods, since the barium sulphate which occurs
on the left of the equation is itself the least soluble of all the
substances likely to be involved. Some other means must
therefore be adopted to bring the barium sulphate into a
soluble form. A common method adopted is to reduce
barium sulphate at a red heat by means of carbon. Reaction
occurs according to the equations —

BaS04 + 4C = BaS + 4CO

Barium sulphate Barium sulphide

BnS04 + 2C = BaS + 2COa

with production of barium sulphide. This barium sulphide
is soluble in acids, and may be easily converted into barium
chloride by the action of hydrochloric acid —

BaS + 2HCI = BaCla + H^S

the sulphuretted hydrogen produced at the same time escaping
as gas.

Silver chloride is another example of a substance which is
insoluble in both water and acids. It also can be brought
into a soluble form indirectly through a process of reduction.
When the insoluble silver chloride is brought into contact
with zinc and acidulated water, the zinc displaces the silver
from its compound in accordance with the list given on page
78, metallic silver and zinc chloride being produced.

2AgCl + Zn = ZnCl2 + 2Ag

The metallic silver obtained in this way may then be con-
verted into silver nitrate by dissolving it in nitric acid, and
from this soluble compound any of the other silver salts may
be prepared. A similar method for obtaining metallic silver
from the insoluble silver sulphide is given in the chapter on
silver.

CHAPTER XVI

OXIDATION AND BEDUGTION

Sometimes we observe what we may call a competition for
oxygen. For example, if we burn a substance like turpentine,
which contains the elements carbon and hydrogen, the flame
produced by the combustion is of a very smoky character,
so that if a cold substance is held near it soot will at once
be deposited. This soot consists of small particles of carbon.
When the turpentine burns, therefore, the hydrogen which
it contains is oxidised in preference to the carbon. Here
we have carbon and hydrogen competing, as it were, for a
limited amount of oxygen within the hot region of the flame,
and in the competition the hydrogen prevails, the whole of
it burning to form water, while only a portion of the carbon
burns to form carbon dioxide, the rest being liberated in
the unburnt state.

There is another sense in which two substances can com-
pete for oxygen. If we take one substance which is already
oxidised, and heat it in contact with another substance which
is not oxidised, but is capable of oxidation, then in some
instances the second substance will take away the oxygen
from the first. An example may be found in the reaction
between lead oxide and hydrogen. If a current of hydrogen
gas is led over heated lead oxide, the following reaction occurs: —

PbO + H2 = Pb + H2O

Lead monoxide Hydrogen Lead Water

Here the hydrogen has taken away the oxygen from the
lead, and is oxidised to water. The lead oxide is said to
have been reduced to metallic lead, and the process is
generally spoken of as one of reduction. Lead oxide may
be similarly reduced by means of carbon. If we heat a
mixture of lead oxide and charcoal, the carbon of the

OXIDATION AND REDUCTION

PbO +

Lead monoxide

Caxbon

2PbO +

charcoal removes the oxygen from the lead oxide and reduces
it to metallic lead, according to the equations —

= Pb + CO

Lead Carbon monoxide

= 2Pb + CO2

Carbon dioxide

The carbon is itself oxidised to carbon monoxide or carbon
dioxide.

Instances of similar actions are given below, the equations
representing the reduction of certain metallic oxides to metals
by means of hydrogen and of carbon —

FeO +

Ferrous oxide

H, =

Hydrc^en

Fe
Iron

+ H2O
Water

SnOa +

Stannic oxide

Carbon

Sn
Tin

+ 2CO

Carbon monoxide

CuO +

Cupric oxide

Hydrc^en

Copper

+ H2O

Water

In what has preceded we have met with another case of
reduction. When red-hot carbon acts on carbon dioxide, the
following action occurs : —

C + CO2

Carbon Carbon dioxide

2C0

Carbon monoxide

Here carbon is oxidised to carbon monoxide, and carbon
dioxide is reduced to the same product.

Carbon monoxide, being capable of still further oxidation,
can reduce certain other oxides. Thus, if a current of carbon
monoxide is led over heated oxide df copper, it removes
oxygen from the copper oxide, which it thus reduces to
metallic copper, and is itself oxidised to carbon dioxide,
the equation being —

CO +

Carbon monoxide

CuO

Copper oxide

CO2 + Cu

Carbon dioxide Copper

One metal is sometimes capable of removing oxygen from
another. If we mix, for example, finely divided aluminium

98 iNORGANIC CHEMIStfeV

with ferric oxide, and start the reaction by heating a portion
of the mixture to a very high temperature, the aluminium
abstracts oxygen from the ferric oxide and reduces it to
metallic iron, in accordance with the equation —

FcgOg + 2AI = 2Fe + AI2O3

Ferric oxide Aluminium Iron Aluminium oxide

So much heat is given out in this reaction that the temperature
of the mass is raised to a very bright white heat, and a mixture
of the above kind has recently come into use as a means of
producing extremely high temperatures.

It will be seen from these instances that reduction and
oxidation go hand-in-hand. When one substance is reduced
another substance is oxidised. If in the action

FeO + H2 = Fe -f HgO

we consider only the ferrous oxide, then we say that the
action is a reducing action, for the ferrous oxide is reduced
to metallic iron. But if we consider the same reaction from
the point of view of the hydrogen, then we must call it an
oxidising action, because the hydrogen combines with oxygen
and is oxidised to water.
Similarly, in the equation —

FegOg + 2AI = AI2O3 + 2Fe

the ferric oxide is reduced to metallic iron, and so the action
is in this sense a reduction; but, at the same time, the
aluminium is oxidised to aluminium oxide, and so the action
must also be considered an oxidation.

A substance which is capable of reducing an oxide is
usually called a reducing agent. Thus, in the above examples
hydrogen, carbon, aluminium, and carbon monoxide act as
reducing agents. The reducing agents which are most
extensively employed on the large scale, chiefly in the
reduction of metals from their oxides, are carbon and carbon
monoxide. These are used in preference to others on
account of their cheapness. Impure carbon can be obtained
in any quantity in the form of coal, or, if greater purity is
desired, of charcoal; and carbon monoxide can be obtained

OXIDATION AND REDUCTION 99

simply by the combustion of these in a defective supply
of air.

The terms oxidation and reduction are sometimes used in
a wider sense than that given to them in the above instances.
Oxidation does not necessarily mean actual union with oxygen
of the element which is said to be oxidised. Thus, it is
customary to speak of sulphuretted hydrogen when it is
converted into sulphur as being oxidised to sulphur. This
action occurs when a solution of sulphuretted hydrogen is
exposed to air, the equation being —

2H2S + 02 = 2H2O + 2S

Now, there is no oxygen in sulphur, and the term oxidation,
when used with respect to the sulphur, here implies that
hydrogen has been removed from it, and not that oxygen has
been added. We very often find the term oxidation used in
this sense to mean removal of hydrogen^ on account of the
fact that oxygen itself can frequently remove hydrogen from
hydrogen compounds, so that addition of oxygen and removal
of hydrogen are in a certain sense actions of the same
type.

Any substance which is capable of adding oxygen to other
substances, or of removing hydrogen from them, is said to be
be an oxidising agent. If we bring chlorine into contact with
sulphuretted hydrogen, sulphur is liberated.

CI2 + H2S = 2HCI + S

The chlorine, by removing hydrogen, has acted as oxygen
acts, and is therefore said to be an oxidising agent. Although,
as we see from the equation, no oxygen is involved at
all, yet the action may be spoken of as one of oxidation
and reduction, the sulphuretted hydrogen being oxidised
to sulphur, and the chlorine reduced to hydrochloric
acid. Reduction, therefore, may not only mean removal
of oxygen, it may also mean addition of hydrogen^ as in the
above instance.

Oxidising and reducing agents are frequently employed in
the laboratory in the liquid or dissolved form. The chief
oxidising agents which are used in this way are nitric acid,
chlorine, hypochlorites, and bromine. Nitric acid, when

loo INORGANIC CHEMISTRY

dilute, does not part with its oxygen very readily at the
ordinary temperature, but when some substances are boiled
with it, it gives up part of its oxygen, and is converted into
less highly oxidised compounds of nitrogen, usually oxides of
nitrogen. Chlorine is an oxidising agent even at the ordinary
temperature, and its solution is employed in the laboratory
under the name of chlorine water, A similar solution
of bromine, called bromine water^ is also employed : it
is a somewhat less powerful oxidising agent than chlorine.
Sodium hypochlorite NaClO, and bleaching powder solu-
tion (which contains calcium hypochlorite Ca(C10)2) are
occasionally used, either with or without addition of acid.
In presence of acid these hypochlorites are more power-
ful oxidising agents than when the solution is alkaline or
neutral.

The following is a ready test for an oxidising substance.
Potassium iodide solution, when oxidised, yields iodine, which
can be easily recognised by its brown colour or by the deep
blue coloration which it imparts to starch solution. Papers
impregnated with potassium iodide and starch are, in conse-
quence, sometimes used in the laboratory as test papers for
oxidising gases. The action of potassium iodide solution
with some of the above oxidising agents may be seen from
the following equations : —

2KI + C12

2KCI + I2

2KI + Brg

■zrr.

2KBr + I3

+ NaClO + H2O

2KOH + NaCl +

2KI + NaClO + H2O = 2KOH + NaCl + I2

Potassium nitrate^ KNO3, and potassium chlorate, KCIO3,

part readily with their oxygen when heated, and so are often
used as oxidising agents in the solid or fused state. Sodium
peroxide^ NagOg, is another substance of the same kind which
has recently come into use.

The following are the chief reducing agents employed in
the laboratory : — Sulphuretted hydrogen, either as a gas or
in solution ; a solution of sulphur dioxide ; a solution of
stannous chloride; and reducing mixtures which yield what
is called nascent hydrogen. Nascent hydrogen has, properly
speaking, no existence. It is supposed to be hydrogen in
the state of formation from the interaction of two substances —

OXIDATION AND REDUCTION loi

Le, hydrogen before it has become ordinary hydrogen gas.
Thus zinc and sulphuric acid, which can act together so as
to produce hydrogen, are capable of jointly reducing certain
substances. Such a reduction is said to be due to nascent
hydrogen, although hydrogen is not necessarily formed at all
if the oxidising agent is present.

The reducing action of such substances can be most easily
shown by means oi potassium permanganate KMn04, which
is a substance with an intense purple colour. When a re-
ducing agent acts upon it in acid solution, it parts with a
p)ortion of its oxygen, and the colour disappears.

When sulphuretted hydrogen acts as a reducing agent, free
sulphur is produced ; when sulphur dioxide acts as a reducing
agent, it is oxidised to sulphuric acid ; stannous chloride is
oxidised to stannic chloride ; and nascent hydrogen is oxidised
to water.

The formation of these substances is shown in the following
equations, in which the symbol (O) indicates oxygen derived
from the substance which is reduced, and not free oxygen
gas:—

Sulphuretted hydrogen

(0)

^^H

,0 + s

Sulphur

SO2 + H2O

Sulphur dioxide

(0)

H2SO4

Sulphuric acid

SnClg + 2HCI
Stannous chloride

+ (0)

1 = SnCU + H2O

Stannic chloride

2(H)

" Nascent hyd^^ogen "

(0)

H2O

Zn + H2SO4 +

(0) =

= H

+ ZnSO^

It is not always easy to find the stage of oxidation of an
element in a given compound by direct reference to the
formula of the compound. If we compare the compounds
SO2 and SO3, it is plain that the sulphur in the latter is more
highly oxidised than in the former, and the same may be said
with respect to the acids corresponding to these anhydrides
H2SO3 and H2SO4. If we are asked, however, to say
whether the sulphur in potassium sulphate Kg SO 4, or in
potassium persulphate KgSgOs, is in the higher stage of

I02 INORGANIC CHEMISTRY

oxidation, it is impossible to answer the question by a direct
inspection of the formulae. The proportion of oxygen to
sulphur is the same in both compounds, but the proportion of
potassium in the two compounds is different. In such a case,
the comparison may most readily be made by first of all
referring these salts to the acids from which they are derived,
and then by subtraction of water, referring these acids to the
corresponding acidic oxides. These acidic oxides contain
only oxygen besides the element considered, and so it is an
easy matter to tell which is the more highly oxidised — that is,
which contains the greater proportion of oxygen. Thus we
have —

KssSO*

H2SO4

SOs or SaOe

K2S2O8

H2S2O8

S2O7

The sulphur in the persulphate is therefore more highly
oxidised than the sulphur in the sulphate.

If we wish to determine the degree of oxidation of a
metallic radical, we proceed in a similar way. Thus, if we ask
whether the iron in ferric chloride FeCls, or in ferrous sulphate
FeS04, is in the higher stage of oxidation, we cannot say off-
hand by inspection of the formulae. It is true that in ferric
chloride no oxygen is present at all, and that in ferrous
sulphate there is a considerable amount of oxygen, but it does
not follow from this that the iron in ferrous sulphate is more
highly oxidised than the iron in ferric chloride. What we
must do is to refer the positive radical to the corresponding
basic oxide. Ferric chloride may be prepared by neutralising
the base ferric hydroxide Fe(0H)3, ferrous sulphate may be
prepared by neutralising the base ferrous hydroxide Fe(0H)2.
If now we subtract the elements of water from these two hydrox-
ides, we are left with the oxide FcaOs in the case of the ferric
chloride, and the oxide FeO in the case of the ferrous sulphate.
There is obviously a greater proportion of oxygen in FcgOs
than in FeO, so we may say that the iron in ferric chloride is
more highly oxidised than the iron in ferrous sulphate,
although in ferric chloride there is no oxygen at all. That
this conclusion is general may be demonstrated by the fact
that any of the oxidising agents just mentioned is capable of
converting a ferrous salt into a ferric salt, and that any of the

OXIDATION AND REDUCTION

103

reducing agents mentioned above is capable of converting a
ferric salt into a ferrous salt.

2FeCl2 +
Ferrous chloride

2HCI +

(O) = 2FeCl8 + H«0

Ferric chloride

2FeCl2
Ferrous chloride

CI.

2FeCl8

Ferric chloride

2FeS04 + H2SO4 +

Ferrous sulphate

FeCls

Ferric chloride

Fe2(S04)3

Ferric sulphate

2FeCl3

Ferric chloride

2FeCl3
Ferric chloride

(O) = Fe2(SOj8 + HgO

Ferric sulphate

FeClg + HCl

Nascent hydrogen Ferrous chloride

SO2 = 2FeS04 + 2H2SO4

Ferrous sulphate

2FeCl2 + 2HCI + S
Ferrous chloride

(H)

nt hj
+ 2H2O +

+ H2S

+ SnCl2

2FeCl2 + SnCl4
Ferrous chloride

In the case of a positive radical, it is easy to tell at once
its stage of oxidation by counting the number of dots attached
to the symbol. The ferrous radical Fe*' has only two dots,
the ferric radical Fe*" has three. The ferric radical is in
a higher stage of oxidation than the ferrous radical, because
each dot represents power of combining with negative
radicals, and these negative radicals may be hydroxyl or
oxygen. The degree of oxidation of a positive radical, there-
fore, corresponds with the number of dots which are attached
to it.

In the case of mercury we have two set of salts, the mercur-
ous salts containing the mercurous radical Hg*, and the
mercuric salts containing the mercuric radical Hg". Now
the mercuric radical has two dots, where the mercurous radical
has only one. We can therefore say at once that the salts
containing the mercuric radical have the mercury in a higher
stage of oxidation than the salts containing the mercurous
radical, and can be formed from the mercurous salts by treat-
ment with oxidising agents.

When a metal passes into the state of a metallic radical,
it assunc^es a charge of electricity whiqh is represented by one

104 INORGANIC CHEMISTRY

or more dots. Thus, when zinc is dissolved in sulphuric acid,
according to the equation —

Zn + H-2S04" = H2 + Zn-SO/

and becomes the positively charged zinc radical, it may be
said to be oxidised. This can be shown by the fact that zinc
and sulphuric acid can act as a reducing mixture, whereas
zinc sulphate, which contains the zinc radical, is entirely
without reducing properties. It may be asked, seeing that
the zinc has been oxidised by the sulphuric acid : What has
been reduced ? The answer is that the hydrogen radical has
been reduced. The hydrogen radical has lost its charge of
electricity and has become uncharged hydrogen gas. In the
above sense, therefore, it has been reduced.

The same thing may be seen by referring the zinc and
hydrogen in zinc sulphate and sulphuric acid back to the
oxides representing their stage of oxidation. If we do this,
we find that the zinc in zinc sulphate corresponds to the oxide
ZnO, and the hydrogen in the sulphuric acid to the oxide
H2O. So far as oxidation and reduction are concerned,
therefore, the solution of zinc in sulphuric acid with pro-
duction of zinc sulphate and hydrogen corresponds to the
action —

Zn + H2O = ZnO + Ha

an action which may be made to take place by passing steam
over heated zinc.

The student will no doubt have observed that so long as he
was concerned merely with the displacement and rearrange-
ment of radicals, the equations to express the reactions were
extremely simple, and could be solved numerically with very
little trouble. On the other hand, equations representing
actions of oxidation and reduction are often somewhat com-
plex, and not to be solved by simple inspection. For example,
if we proceeded by a process of trial and error to solve the
equation —

?Cu + PHNOs = ?Cu(N03)2 + ?N0 + ? HgO

we might make a great many trials before we found the pro-
portions 3CU and 8HNO3, which satisfy this equation numeric-
ally. It is always easy, however, by systematic procedure to

OXIDATION AND REDUCTION 105

arrive at the correct numerical solution even with equations
much more complicated than that given above. The actual
method adopted may be varied to suit particular cases, but the
essential feature is to split up the total reaction into a series
of simpler reactions, the equations for which may be arrived
at by inspection, and then to combine these so as to give the
equation for the complex reaction.

One part of the action of nitric acid on copper consists in
oxidation of the copper, and reduction of the nitric acid. We
may begin, then, by writing an equation to express the reduc-
tion of the nitric acid to nitric oxide, and so ascertain the pro-
portion of oxygen available for oxidising the copper.

The equation to express the decomposition of nitric acid
into water, nitric oxide, and oxygen is as follows : —

PHNOg = PHaO + ?N0 + ?0

It is evident that the solution of this equation is

2HNO3 = H2O + 2NO + 3O

The quantity of oxygen represented by 3O is not given off as
such, but goes to oxidise copper to the stage of oxidation in
which it is contained in copper nitrate. Referring back to the
basic oxide from which copper nitrate is derived, we find that
this stage of oxidation is represented by the formula CuO.
Each Cu which passes into CuO requires O, therefore to use
up all the oxygen available from the nitric acid we must take
3CU, and then arrive at the equation —

3CU + 30 = 3CUO

Now, all this copper oxide must be converted into copper
nitrate, which can easily be done by acting on the copper oxide
with nitric acid. No oxidation or reduction being involved
in this process, we arrive at once at the solution —

3CUO + 6HNO3 = 3Cu(N03)2 + 3H2O

If we now add these three equations together, we get —

3CU + 3O + 3CuO + 8HNO3 =
3Cu(N03)2 + 3CUO + 4H2O + 2NO + 3O

io6 INORGANIC CHEMISTRY

or, if we strike out terms which are common to both sides,
3Cu + 8HNO3 = 3Cu(N08)2 + 4H2O + 2NO

This last equation is the correct equation for the action of
nitric acid on copper with formation of the substances copper
nitrate, nitric oxide, and water.

Suppose we have to solve the equation —

?Zn + PHNOg = ?Zn(N03)2 + ? NgO + ? HgO

we may proceed in exactly the same way as before. First of
all, we find out how much oxygen is available for oxidation
when nitric acid is reduced to nitrous oxide. This quantity is
given by the equation —

2HNO8 = H2O + N2O + 4O

Now, zinc in zinc nitrate is in the stage of oxidation corre-
sponding to the oxide ZnO. To use up all the oxygen avail-
able from the nitric acid we must therefore write —

4Zn + 40 = 4ZnO

and to convert the amount of zinc oxide thus obtained into
zinc nitrate we must write —

4ZnO -f- 8HNO3 =4Zn(N03)2 + 4H2O

Adding these three equations together right and left, and
striking out members common to both sides, we get the
solution —

4Zn + loHNOs = 4Zn(N03)2 + 5H2O + NgO

for the action of zinc on nitric acid, with formation of nitrous
oxide as the reduction product of nitric acid.

Under certain conditions nitric acid acts upon zinc with
production of ammonium nitrate.

?Zn + PHN03 = ?Zn(N03)2 + ? NH4NO3 + ? H^O

The student might find this equation somewhat complicated
to solve, but having arrived at the solution of the equation
which represents the formation of nitrous oxide, he can easily
deduce the equation representing the formation of ammonium

OXIDATION AND REDUCTION 107

nitrate by finding out in what way nitrous oxide and ammonium
nitrate are related to each other. If we write the formula of
ammonium nitrate N2H4O3, and deduct from this the formula
of nitrous oxide N2O, we find there remains as residue H4O2,
or 2H2O. We can therefore write the equation —

N2O + 2H2O = NH4NO8

This equation represents a purely imaginary reaction (the
reversed equation, however, being true), but since this reaction
cuts out in the final result, it is permissible to make use of
it. If we now add the two equations —

4Zn + loHNOs = 4Zn(N03)2 + N2O + 5H2O
N2O + 2H2O = NH4NO8

and cut out the terms common to both sides, we arrive at the
equation —

4Zn + 10HNO3 = 4Zn(N03)2 + NH4NO3 + 3H2O

The student is advised to practise this mode of dissecting
complicated reactions into simpler reactions (as far as possible
into those with which he is familiar), and then building up the
complicated equation by means of the equations for these
simpler reactions.

CHAPTER XVII

THE GENEBAL LAWS FOR GASES

If we wish to ascertain the quantity of a gas, we usually find it
more convenient to measure its volume rather than to weigh it,
as the weighing of gases is a troublesome operation. But it
must be borne in mind, that while the weight of a given quan-
tity of gas remains unchanged under all conditions, the volume
varies very much according to the conditions under which we
measure it. In the first place, the volume of a gas varies with
the pressure upon the gas. The variation, however, takes place
according to a fixed law, which is not only always the same for
a given gas, but is the same for all gases. This law is known as
Boyle's law, and states that the volume of a given quantity of
gas is inversely as the pressure on the gas, if the temperature
remains constant. Thus, if we reduce the pressure on a gas to
one-half, the gas will double its volume ; if we double the
pressure on a gas, the gas will be compressed to half its volume ;
if we quadruple the pressure on a gas, it will be compressed to
one-fourth of its original volume ; and so on. This law can be
expressed otherwise by saying that the product of the pressure
and volume of a given quantity of gas is constant at constant
temperature. If the pressure on a gas be represented by/,
and the volume by v, then we have the equation —

pv = constant,

provided that the temperature always remains the same.

Not only are all gases affected equally by change of pressure,
their volume is also affected equally by a given change in the
temperature. The law regulating the change of volume of
gases caused by change of temperature is known as Gkty
Lussac's law, and may be stated as follows : —

The volume of a given quantity of gas is directly proportional
to its absolute temperature, provided the pressure remains
constant. If we use the centigrade divisions, the absolute

io8

THE GENERAL LAWS FOR GASES

109

temperature of a substance is equal to its centigrade tempera-
ture plus 273. Thus the absolute temperature corresponding

No pn»aure-

laU

7 atm.l

"'1 aim.

2 aim.

— / aim.

Fig. 19. — Diagram illustrating Boyle's Law.

B represents a mercurial barometer showing the average atmospheric pressure.
/, //, and /// show the volumes occupied by the same amount of gas at pressures
of I, a and ^ atmospheres respectively.

to io°C. is 10 + 273=283. The volume of a given quantity
of gas will be doubled if we heat from o°C. to 273X., for the

no INORGANIC CHEMISTRY

absolute temperatures corresponding to these are + 273 and
273 + 273 : the absolute temperature has been doubled, and
consequently the volume of the gas is doubled.

Suppose now that the pressure and the temperature of a gas
are both changed. By combining the above laws we arrive at
the following expression : — If /^, v^, T^, are the pressure,
volume, and absolute temperature of a gas under one set of
conditions; /„ v^, T„ the corresponding magnitudes under
another set of conditions, then

If any five of these six magnitudes are known, we can calculate
the remaining magnitude. Thus we are in a position to solve
a problem like the following: — "If a gas is heated from lo^C.
to 2o°C., by how much will the pressure on it have to be in-
creased in order to bring it back to its original volume ? "

Substituting the numerical values for the temperature in the
above equation, and putting as the conditions of the problem
require v^ = v^y then —

283 293

and the increase of pressure is -^^ of the original value.

Knowing these simple laws for gases, it is possible to read
the volumes of gases under any conditions which may be found
convenient, and calculate from that volume and these condi-
tions the volume which the gas would occupy under what we
call normal conditions. Normal pressure we take to mean the
average pressure of the atmosphere, which is equal to the
pressure of a column of mercury 760 mm. high. The normal
temperature is the temperature of melting ice — namely, o°C.
It is convenient to have these standard conditions for measur-
ing the volumes of gases, for it enables us easily to pass from
the volumes of gases to their weights.

On account of the gas laws holding for all gases — that is, on
account of all gases being affected equally by changes in tem-
perature and pressure, we can compare the volumes of gases with
each other, not only under the standard conditions but under

THE GENERAL LAWS FOR GASES iii

any conditions, provided they are the same for the gases com-
pared. Thus, if under one set of conditions the volume of a
certain quantity of one gas is equal to twice the volume of a
certain quantity of another gas, then, no matter how the condi-
tions are changed, the volume of the first gas will always be
double the volume of the second gas, if the volumes of the two
gases are measured under conditions which are the same for both.

Solubility of Gases. — When a gas is only moderately soluble
in a liquid, the amount of it dissolved by a given quantity of
the liquid depends upon the pressure, and the solubility is
generally stated as the number of volumes of the gas dis-
solved by one volume of the liquid. The manner in which
a quantity of gas dissolved varies with the pressure is very
simple, and is known as Henry's law. Af a given temperature
the weight of a gas dissolved by a given bulk of liquid is propor-
tional to the pressure of the gas.

Thus, at o'C. water dissolves 4 per cent, of its own volume
of oxygen at one atmosphere pressure; at two atmospheres
pressure it would dissolve twice as much by weight ; at three
atmospheres three times as much ; and so on. It must not
be supposed, however, that at two atmospheres pressure the
water will dissolve 8 per cent, of its own volume of oxygen,
for at two atmospheres pressure each volume of oxygen gas
will contain twice the weight of oxygen that it did at one
atmosphere pressure. Thus, doubling the pressure doubles
the weight of gas dissolved, but it halves the volume which
the gas occupies before it is dissolved. If we then consider
volumes, we may state Henry's law in the form that the volume
of a gas dissolved by a given volume of liquid is independent of
the pressure. We can thus say that at o'*C. water will always
dissolve 4 per cent, of its volume of oxygen, no matter what
the pressure of the oxygen is, although, of course, the actual
weight of oxygen dissolved varies directly with the pressure.

In giving the solubility of a gas in water it is customary
to reduce the dissolved volume, measured under the condi-
tions of the experiment, to the volume which it w^ould occupy
at o*. This reduced volume is usually spoken 6f as the
absorption coefficient of the gas.

The solubility of different gases in water is very variable,
as the following table shows. The numbers given are the

112 INORGANIC CHEMISTRV

volumes of gas dissolved by one volume of water at o° and
760 mm.: —

Ammonia, NHg . 1050 volumes.

Hydrochloric acid, HCl . 505 „

Sulphur dioxide, SO,

Sulphuretted hydrogen, HaS

Carbon dioxide, COj

Argon, A

Oxygen, Oj
Nitrogen, Nj
Hydrogen, H^

Hi/drojen

,04

The solubility of gases almost invari-
ably falls off as the temperature rises,
and gases may usually be expelled
from water by boiling. Gases which,
like ammonia and hydrochloric acid,
are excessively soluble in water, do
not obey Henry's law.

position of Water.

sulphuri

Reacting VolnmeB of Oases. —

When we measure the volumes of
gases which take part in chemical
actions, we find that they are related
in a very simple way to one another,
provided that the volumes are all
measured at the same temperature
and pressure. Thus, when water is
decomposed by electrolysis, two vol-
umes of hydrogen are produced at
the kathode for each volume of
oxygen produced at the anode. If
we reverse this action, we find that

^ _ _^ exactly these proportions of oxygen

omposed hy and hydrogen unite to form water.

^._^ ,d not only so, but that the volume

niaybeaeenLniheaboveBppBraiui Qf water-vapour produced is exactly

orQire="p"'Su«ii«ihe™n^f! equal to the volume of hydrc^en

J^itTb^'iKKd'' ThTI^'ier'Ss burned. Again, hydrogen and chlo-

fotwdbyiheaccuinuiaiinggasto rine Unite in equaJ volumes to pro-

nie loio the bulb c. Amcg hydrocbloric acid gas, the vol-
ume of which is exactly equal to the sum of the volumes

THE GENERAL LAWS FOR GASES

of the hydrogen and chlorine ; and when carbon or
sulphur is burned in oxygen, the volume of the carbon
dioxide or sulphur dioxide pro-
duced is exactly equal to the
volume of oxygen which has

This simplicity in the propor-
tions by volume of gases entering
into chemical action was first
noticed by Gay Lussac, and is
usually called Oay Lnesac's law
of Tolnmes.

Now, if we write the equations
for the actions which have just
been mentioned, noting the re-
acting volumes alongside, we
have —

2H„

jH^O

I vol.

2 vols.

3HC1

I vol.

3 vols.

f°f.

"t;

21. — Volumetric Composi-
1 of CaiboQ Dioxide or Sul-

SOa

ph«

,r Dioxide.

I vol.

quatior

isitii

1 obvious

A pi.

«:e of chaicoa! or sulphur placed
spoon S may be made lo ignite
ojiynn wiih which the crosed

that a very simple relation exists t'^^n'^^ti'Jl"J'"u^n,"'"fi!^
between the volumes of the gases the combunion ii campieied and the
and their formula- weights as we *?* ™i^bii°fo'ind'i£i't'hl""iu™ of
have written them. If, for each the nas, ^ indicated by ih= levei. ot
formula of gas in the weight w rtJ^o^''unX!lJgi!'" """' '"'''
equation, we write volume, we

arrive at the proportions by volume in which the gases
actually enter into the chemical action. As a matter of
fact, our formula-weights are chosen so that this very simple
relation between the formula and volumes of gases may
appear.

Instead of having written the formulse of hydrogen and
chlorine, H, and Clj, we might have^adopted the simpler

114 INORGANIC CHEMISTRY

formulae H and CI. The equation for the formation of
hydrochloric acid gas would have then become —

H + CI = HCl

This equation expresses perfectly the weight and composition
of the substances involved in the action, but we should lose
by adopting these formulae the simple relation between
formula-weights and gaseous volumes, for the formulae H and
Q would each stand for one volurne, and the formula HCl for
two volumes ; so that if we substituted volume for formula we
should no longer arrive at the correct reacting proportions by
volume. Again, if we wrote —

H2 + O = H2O

we should have the correct proportions by weight, but would
not obtain the correct proportions by volume on substituting
volume for formula.

The formulae of gases which are selected in accordance
with the above simple relation — namely, that the formula
should all represent equal volumes of the different gases — are
called molecular formuls, and the weights of the gases which
they represent are called the molecular formula-weights, or,
shortly, the molecular weights of the gases.

The molecular weights of different gases are thus propor-
tional to the weights of equal volumes of these gases, all
volumes being of course measured at the same temperature
and pressure. In other words, the molecular weights of gases are
proportional to their relative densities. This statement is known
as Avogadro's principle, and is used for practically fixing the
molecular weights of gases. If we fix the temperature and
pressure at which all gases are supposed to be measured at
the normal values of zero and 760 mm. (N.T.P.), we may give
the following numerical expression to Avogadro's principle : —

The molecular weight of a gas is equal to the number of
grams of the gas which would occupy 22*4 litres at O"" and
760 mm {see pp. 166-7).

The molecular weight of gaseous substances expressed in
grams is often spoken of as the gram molecular weight, and
the volume of 22.4 litres which this gram molecular weight
of a gas would occupy at 0° and 760 mm. is frequently
called the gram molecular yolnme of the gas at N.T.P,

THE GENERAL LAWS FOR GASES 115

The following is an example of an actual calculation of a
molecular weight: — It is found that 21.5 litres of oxygen
weigh 28.1 grams at i9°C. and 744 mm., what is the mole-
cular weight of oxygen ? The problem actually to be solved
is this : If 28.1 grams of a gas at iq^'C. and 744 mm. occupy
21.5 litres, what weight of it will occupy 22.4 litres at oX. and
760 mm.? We first reduce the volume under the given
conditions to N.T.P. as follows : —

744x21.4 _ 76o:r

273+19 273

X = 19.6

Since 28.1 grams occupy 19.6 litres at N.T.P., what weight
will occupy 22.4 litres?

22.4 X 28.1

— =^—7 — = 32.0.

19.6

The molecular weight of oxygen is thus 32, and since the
symbol O stands for 16, we have O2 for the molecular
formula of oxygen.

In order to make such a calculation of molecular weights, it
is not necessary that the substances considered should be gases
at the ordinary temperature. We can find the molecular
weight of water, which is a liquid at the ordinary temperature,
by determining the weight of a given volume of its vapour
at any temperature above the boiling point, and then reducing
that volume to normal conditions, although it is impossible
actually to have gaseous water existing at 0° under 760 mm.
pressure. Such a determination is usually called a measure-
ment of the yaponr density of a substance.

When dealing with gases, chemists invariably use mole-
cular formulae in their equations, for then they are enabled
at once to pass* from weights to volumes, or vice versd, by
means of the rule given above.

It must be remembered, however, in making this passage,
that the volume of 22.4 litres holds good for 0° and 760 mm.
only. If the actual volumes considered are measured under
other conditions of temperature and pressure, they must be
reduced to the normal conditions before the passage from
volumes to weights can be made.

Suppose we have to solve the following question : — What

ii6 INORGANIC CHEMISTRY

volume of carbon dioxide at N.T.P. can be produced by
heating 20 grams of calcium carbonate to redness ? We first of
all write the equation for the action, using molecular formulae
for the gases, and then note alongside the substances whose
weights are required, the weight in grams represented by the
formulae, and alongside the gaseous substances 22.4 litres for
each molecular formula of the gases, thus —

CaCOa = CaO + CO2

100 g [56 g] [44 g]

— — 22.4 litres

We at once see from this equation that 100 grams of calcium
carbonate give 22.4 litres of carbon dioxide at N.T.P.,
and therefore that 20 grams give 4.48 litres under these
conditions.

Again, we might be required to solve a question like the
following : — How much calcium carbonate must be heated in
order to give one litre of carbon dioxide measured at 2o°C.
and 770 mm.? We first of all reduce this volume to N.T.P.
as follows : —

770 X I = 760^

293 273

X = 0.94

From the above equation 22.4 1 are yielded by loog, so that
0.94 1 will be yielded by

100 X 0.04

— = 4.2 g.

22.4

For purposes of calculation the student will do well at
first always to write down the equations for the actions
involved, each formula being accompanied by the weight in
grams which it expresses, and each molecular formula of a
gas by the corresponding volume in litres. Thus, for
example, he should write —

H2 + Clg = 2HCI

2g 71 g 73 g

22.4 1 22.4 1 2 X 22.4 1 at N.T.P.

This will enable him to solve any question with regard to the
weights and volumes of the gases entering into this reaction.

CHAPTER XVIII

GASEOUS MIXTURES

All gases are capable of mixing with each other naturally — i.e.
of diffusing into each other. If a bottle containing ammonia,
or any other gas which is recognisable by the smell, is carefully
opened, then even although the air is free from mechanical
disturbance the smell of the ammonia will very soon be per-
ceptible at a considerable distance from the bottle, showing
that the ammonia gas must have diffused into the gases of the
atmosphere.

The process of gaseous diffusion may also be rendered
visible by choosing a coloured gas and a colourless gas for
the experiment. If a small bulb containing liquid bromine is
crushed at the bottom of a tall cylinder, the liquid bromine
on escaping is partially converted into bromine vapour — i.e.
bromine gas, which is easily recognised by its dark -brown
colour. If the cylinder is left to itself, the bromine vapour
will be found to rise gradually, the tint throughout the cylinder
not becoming uniform until several hours have elapsed. Now,
bromine vapour is bulk for bulk over five times heavier than
air, and the action of gravity would tend to keep the heavy
vapour at the bottom of the cylinder, yet the bromine
moves upwards at a considerable rate. If the cylinder is
originally filled with hydrogen instead of with air, the rate of
diffusion can be seen to be very much greater than in the
former instance, notwithstanding the fact that bromine is
eighty times heavier than an equal bulk of hydrogen. The
gravitational action against diffusion in this second case is
much greater than in the first, yet the diffusion proceeds at a
greater rate. We have, however, been looking at the process
of diffusion only from the point of view of the coloured bromine
gas, but this gas alone is not responsible for the whole of the
diffusion or mixing process. The air in one case and the

117

ii8

INORGANIC CHEMISTRY

hydrogen in the other really play the principal part in the
mixing. It is found that the lighter a gas is the faster it
•diffuses. Now hydrogen is 14 J times lighter than air. It
diffuses, therefore, more rapidly than air, and the process of

mixing of hydrogen and bromine
is consequently more rapid than
the mixing of air and bromine.

The relation between the den-
sity of a gas and the rate at which
it diffuses may be stated quite de-
finitely. The speed of diffusion of
a gas has been found to be inversely
proportional to the square root of
its density. Oxygen is sixteen
times as heavy as hydrogen.
These gases will therefore have
speeds of diffusion proportional
to—

Fig. 22. — Diffusion of Gases.

^16 : ^\ or

1
T

The bulb C, of porous earthenware,
has a long glass tube fastened into its
neck by means of a cork. The end of
this tube dips under water in the
beaker A , and over the bulb a beaker
B, filled with hydrogen, is inverted.
Since hydrogen diffuses much faster
than air, it enters C through the pores
more rapidly than the air can pass
out through them. Gas thus ac-
cumulates mside the bulb and escapes
through the water in A.

Collection of Gases. — These
considerations are of importance
when we come to deal with
practical methods for collecting
and manipulating gases. If a
gas is heavier than air, it may
be collected by downward dis-
placement in the manner shown
in the figure. The gas is de-
livered at the bottom of the jar, and since diffusion is a
comparatively slow process, it there forms a layer heavier
than the air, the surface of this layer gradually rising as
more gas is delivered, and displacing the air with which the
jar was originally filled. If the heavy gas is delivered slowly
there is little mechanical disturbance, so that the gas and air
can only mix by diffusion. It is possible in this way, then,
to collect a sample of a heavy gas which will contain very little
air.

If the heavy gas were delivered at the top of the cylinder

GASEOUS MIXTURES

ii9

Air]

Heavy
Gas

light
6aa

Air

it would, in virtue of its greater weight, tend to fall to the
bottom of the cylinder; we should therefore have a current
of the heavy gas proceeding downwards and the lighter air
coming upwards. This would involve considerable disturbance,
so that the two gases would mingle not only by diffusion but
by actual mechanical mixing.

When the gas which it is desired to collect is lighter than
air, it may be collected by upward displacement^ as shown in
the figure. The jar in
which the gas is to be col-
lected is inverted so that
the open mouth is down-
wards. The delivery tube
ends at the upper portion
of the jar. The light gas as
it is delivered forms a layer
at the upper end of the jar,
* which gradually increases in
size, forcing the heavier air
downwards. There arc thus
no conflicting currents of
the two gases, and, con-
sequently, no mechanical
mixing. Any process of
mixing that goes on must
be due to the natural diffu-
sion of the gases.

It has just been said that
air is about 14.5 times as
heavy as hydrogen — ue,
about 29 grams of air can be contained in 22.4 litres at
N.T.P. Any gas, then, whose molecular weight is greater
than 29, is heavier than air, and can be collected by down-
ward displacement; any gas whose molecular weight is less
than 29, is lighter than air, and can be collected by upward
displacement. Examples of the former class are CO 2, HCl,
N2O, O2, the corresponding weights being 44, 36.5, 44* 32 :
examples of the second class are Hg, NH3, CH4, with the
weights, 2, 17, 16 respectively.

In the above methods, the gases which it is desired to
collect are made to displace the air with which the vessels are

Fig. 23. — Collec-
tion by Down-
ward Displace-
ment.

Fig. 24. — Collec-
tion by Upward
Displacement.

The dotted lines in the figures show the
surface of separation of the gases.

120 INORGANIC CHEMISTRY

originally filled. Another very common method of collecting
gases is to collect them over a liquid. This method resembles
the method of collection by upward displacement; only the
vessel is originally filled, not with air, but with a suitable
liquid. The resemblance of collection by upward displace-
ment depends, of course, on the fact that all gases are much
lighter than liquids. In selecting a liquid for the purpose,
it is evident that one must be chosen in which the gas is
insoluble, or only slightly soluble, for otherwise the gas which
it is desired to collect would dissolve in the liquid, and be lost
as gas. The liquid, usually water, is contained in a compara-
tively shallow vessel, in which the collecting jars, also filled with
the liquid, are immersed mouth downwards. The end of the
tube which delivers the gas is brought immediately beneath
the mouth of the jar in which the gas is to be collected. The
gas bubbles up through the liquid in the jar, displacing it and
forcing it downwards, and the process can be continued till all
the liquid is displaced (see fig. 26, p. 127).

Partial Volnine and Partial Pressure. — When we are
dealing with a mixture of gases we can consider the mixture
from two points of view. In the first place, it should
be noted that gases, when mixed, do not (for ordinary
purposes) influence each other's pressure or each other's
volume. We can consider, therefore, that the total volume
of the gas mixture is the sum of the volumes of the separate
gases it contains. Thus we say that, roughly speaking, air is
a mixture containing one volume of oxygen, and four volumes
of nitrogen. This means that if the air were separated into
its components, the nitrogen would occupy four-fifths of the
original volume, and the oxygen would occupy one-fifth, pro-
vided all measurements were made under the same conditions
of temperature and pressure. These volumes may be called
the partial volumes of the components, so that we can say
that the total volume of a gas is equal to the sum of the partial
volumes of its components.

The other way of considering a mixture of two gases is to
look at it from the point of view of their pressures. We can
suppose the mixture to be separated into its two component
gases, each component occupying a volume equal to the
original volume of the mixture. Thus, we can suppose the

GASEOUS MIXTURES 121

nitrogen and the oxygen of air to be separated from each other,
and each to occupy the whole volume occupied by the air
from which they were obtained. Now, if we consider that at
the pressure of the original air oxygen occupies only one- fifth
of the original volume of the air, then if it is expanded so as
to occupy a volume equal to the whole volume occupied by
the air, its pressure, by Boyle's law, will be reduced to one-
fifth of the original pressure. Similarly, if the nitrogen is
expanded so as to occupy, not four-fifths of the original
volume, but a volume equal to the total original volume, then
its pressure will be reduced to four-fifths of the original pres-
sure. We can therefore say that the total pressure of the air
is composed of the pressure of oxygen, which is equal to one-
fifth of the total, and of the pressure of nitrogen, which is
equal to four-fifths of the total. These pressures are called
the partial pressures of the components, and we may state in
general that the total pressure of a mixture of gases is equal
to the sum of the partial pressures of its components.

When we are dealing with the solubility of a mixture of
gases, we find that each gas dissolves independently of the
presence of the other gases, and we may state this in the
following form, which is usually known as Dalton's law of
partial pressures : — Each gas in a mixture dissolves according
to its own partial pressure. For example, we may investigate
the solubility of air in water, supposing that it consists of a
mixture of oxygen and nitrogen, which for rough purposes of
calculation we may represent by the proportions of \ to ^,
The absorption coefficient of oxygen in water at o°C. is 0.04 ;
the absorption coefficient of nitrogen at the same temperature
is 0.02. Now, if we suppose the pressure of the atmosphere
to be at its average value, then the quantity of oxygen which
will be dissolved by one volume of water will amount to
0.04 X ^ = 0.008, this volume being measured at one
atmosphere pressure. Similarly, a quantity of nitrogen dis-
solved by one volume of water will be 0.02 x |^ = 0.016, this
volume being again measured at one atmosphere pressure.
Measuring by volumes, then, the quantity of oxygen dissolved
by the given quantity of water is equal to ^ the volume of
the dissolved nitrogen, instead of being \ of the volume, as
it was in the original mixture.

CHAPTER XIX

THE ATMOSPHERE

It has already been stated that air is a mixture consisting
chiefly of nitrogen and oxygen. In view of what we have
learned concerning the behaviour of a mixture of gases, we
can give the composition of the air either in terms of the
partial volumes, or of the partial pressures. The first method
is that usually adopted in giving the composition of a mixture
of gases ; but, as we have seen, it is sometimes convenient to
adopt the second method in considering some of the properties
of such mixtures. The average composition by volume of
what we term purified air is—

Nitrogen, N2 78.2 volumes per cent.
Oxygen, Oj 21.0 „
Argon, A 0.8

TOCO

The partial pressures of these gases are represented by the
same numbers. If we wish to know the composition of the
air by weight, we can easily calculate it* from the molecular
weights of the component gases, and the proportions in which
these gases occur in the mixture. The calculation is per-
formed as follows: —

78.2x28 = 2190 = 74.8 per cent.
21.0x32 = 707 = 24.1 „

0.8x40 = 32 = I.I „

2929 loo.o

The argon in the atmosphere was only discovered in 1894,
although the quantity of the gas in the air must be estimated
in billions of tons. The reason why chemists were so long
in detecting its presence is that, like nitrogen, it is a very

122

THE ATMOSPHERE 123

inert gas, and takes no part in those chemical actions in which
atmospheric air is one of the reacting substances. It was lost
sight of in the very much larger quantity of nitrogen, and was
only detected by the slightly greater weight of " atmospheric
nitrogen " (that is, the mixture of nitrogen and argon), when
compared with the weight of pure nitrogen prepared from a
chemical compound.

That the composition of the air should be practically
constant is no proof of the chemical union of its components,
for, as we have seen, gases mix with each other very readily,
and the atmosphere is being perpetually disturbed by air
currents in the form of wind.

It may be shown in many different ways that the nitrogen
and oxygen in the air are not in chemical combination. In the
first place, the quantities of nitrogen and oxygen in the air
are not related in any simple way to the combining weights
of these elements. In the second place, if pure oxygen and
nitrogen in ihe requisite proportions are simply mixed, no
sign of any chemical action can be detected, and yet the
mixture has all the properties of ordinary purified air. That
air is a mixture of oxygen and nitrogen can also be proved by
considering the manner in which it dissolves in water.

We have seen that if air consists of a mixture of oxygen and
nitrogen, in the proportions of ^ to ^, these proportions should
be altered when the gases dissolve in water, owing to their
different solubilities, and that, if we recovered the dissolved
gases from the water, their proportions should then be ^
oxygen to f nitrogen (p. 121). The dissolved air can readily be
expelled by boiling the water, for at the boiling point neither
nitrogen nor oxygen is appreciably soluble in water. The
gases thus evolved can be collected, and the composition
of the gaseous mixture can be ascertained. When this
experiment is performed, it is found that the dissolved air
is much richer in oxygen than air before being dissolved
in water, and that the proportions are very nearly those
required by the above calculation. If the air had been
a compound of oxygen and nitrogen, it would have been a
single gas having its own single absorption coefficient, and
the composition of the gas expelled from the water by
boiling would have been exactly the same as the composition
of the gas before it was dissolved.

H4 INORGANIC CHEMISTRY

If we wish to ascertain experimentally the quantity of oxygen
in the air, one of the following methods may be employed : —
Oxygen being the active constituent of the atmosphere, may
be made to combine with many other substances, the nitrogen
and the ai^on, which together constitute what is still frequently
called "atmospheric nitrogen," being
unaffected. It is easy, then, to choose
some substance which will combine with
the oxygen and leave the atmospheric
nitrogen behind. One such substance
is phosphorus. If we leave a stick of
yellow phosphorus exposed to moi^t air,
the phosphorus combines slowly with the
oxygen at the ordinary temperature, and
after a p>eriod of about an hour, all the
oxygen will have been removed and con-
verted into oxides of phosphorus, which
dissolve in water. By measuring the
original volume of the mixture, and then
measuring the volume of the atmos-
pheric nitrogen which remains after the
oxygen has been removed, making due
allowance for temperature, pressure, and
the presence of moisture, the composi-
_ tion of the mixture by volume is at once

F,g. IS -P™pon.on of ascertained.

OxvEcn in All. . , , ,- ,

A more exact method may be applied
inla™u^''i'X"3'if«' to the dry mixture. We have seen that
po^ov=r™i«-toihe»aioD copper when heated in oxygen com-
phAJ'^'fusoTioThe'end bincs With the Oxygen to form copper
»ciroii7"compittt ihe'TOium* 0^'de. In a specially constructed ap-
afrBiduiiguiimcuured. paratus, a measured volume of dry air
may be brought into contact with
copper wire which is heated to redness by means of an electric
current passing through it The copper, when it is heated,
takes up all the oxygen in the air, and leaves the atmos-
pheric nitrogen behind. A measurement of the volume of
this atmospheric nitrogen gives, by subtraction from the
original volume, the volume of oxygen which has been re-
moved, and thus the volumetric composition of the purified
dry air is ascertained.

THE ATMOSPHERE 125

Atmospheric air, besides containing these permanent consti-
tuents, also contains considerable quantities of other gases.
The gas which exists in greatest quantity, although this
quantity is liable to great variation, is water-vapour. It can
be easily understood that water-vapour occurs to a large
extent in the atmosphere, if we consider that nearly three-
fourths of the earth's surface is covered by water, with which
the atmosphere is in constant contact. This water evaporates,
and consequently we find that the air is always more or less
moist.

At each temperature air is capable of taking up a definite
amount of moisture. Sometimes the full amount of moisture
is found in the air, and the air is then said to be saturated with
moisture. This occurs in fog or mist, or during a heavy rain.
As a rule, however, the air is only about two-thirds saturated,
and the average amount of moistiu-e contained in the air at
the mean temperature of the atmosphere is about 1.4 volumes
per cent.

Carbon dioxide is also invariably present in the atmosphere,
and although the proportions in which it is found do not vary
so much as in the case of the atmospheric moisture, yet they
are not so constant as the proportion of the permanent gases
nitrogen, oxygen, and argon. In country air, and in the air
over the ocean, the proportion of carbon dioxide is about 3
volumes in 10,000 volumes of air. This proportion does not
fluctuate very much in such regions, but in towns the propor-
tion of carbon dioxide in the air is considerably higher, the air
in the streets containing usually about 4 volumes of carbon
dioxide in 10,000. The larger amount of carbon dioxide in
towns arises from the respiration of living beings in these
places, and from the amount of carbon consumed as fuel. In
inhabited rooms, especially in ill-ventilated rooms in which
gas is burnt, the proportion of carbon dioxide frequently
rises to 10 volumes in 10,000, and in crowded apartments
occasionally reaches as much as 50 volumes in 10,000.
Carbon dioxide is not in itself a very poisonous gas, but its
presence in the air in excessive quantity indicates undue
contamination of the air.

The quantity of carbon dioxide present in air may be
estimated in various ways. The methods, however, chiefly
depend upon the absorption of the carbon dioxide by a

126 INORGANIC CHEMISTRY

liquid such as lime water, or baryta water. The absorption
takes place according to the equations —

Ca(0H)2 + CO2 = CaCOs + HgO
Ba(0H)2 + CO2 = BaCOa + H^O

the calcium or barium carbonate formed being insoluble. If
we know the strength and amount of the calcium or barium
hydroxide solution originally taken, and ascertain its con-
centration after it has been brought into contact with a measured
volume of air, then by finding what the diminution in strength
is, we can tell how much carbon dioxide the air originally
contained.

Air can be freed from moisture and from carbon dioxide by
passing it in succession through tubes containing concen-
trated sulphuric acid, to remove water, and soda-lime, to
remove carbon dioxide. When we speak oi purified air^ we
usually mean air which has been treated in this fashion.

Recent observations have shown that besides the gases
previously mentioned, ordinary atmospheric air contains a
great many other gases which are present only in very small
quantity. Thus, in the air have been found small amounts
of ammonia, hydrocarbons, hydrogen, helium, and other gases
of a similar nature.

CHAPTER XX

OXYGEN

Since oxygen is the active constituent of the air, it is
somewhat difficult to separate it as such from the atmosphere.
The plan which must be adopted to get pure oxygen from
the air in quantity is first of all to make the atmospheric
oxygen combine with some other substance, and then, by
suitably treating the product of the union, recover the oxygen
from this.

One method, although not a convenient one, for doing this,
we have already met with. If we heat mercury in the air
to its boiling point, it will very slowly combine with a little
oxygen to produce mercuric oxide, according to the equation —

2Hg + 02 = 2HgO

This mercuric oxide may then be heated to a somewhat higher
temperature, when it splits up again into metallic mercury

Fig. 26. — Preparation ol Oxygen from Mercuric Oxide.
The oxide is heated in a test-tube and the oxygen gas given off collected over water.

and oxygen, which by means of suitable apparatus can be
collected in the pure state. The mercury recovered from
the mercuric oxide can once more be made to unite with

127

128 INORGANIC CHEMISTRY

oxygen in the air, and the oxide formed again decomposed.
Thus a very small quantity of mercury, by repeated use, can
be made to remove from the air a comparatively large
quantity of oxygen, which can be obtained in the pure state.
This method, although of historical importance, would be
so excessively tedious and troublesome that it is practically
never adopted, except for purposes of illustration. Oxygen,
however, may be obtained nearly pure from the atmosphere
by means of two reactions, which are quite analogous to
those just considered. If barium oxide is heated to a high
temperature in presence of air, it combines with the oxygen
and forms barium dioxide.

2BaO + 02 = 2Ba02

Barium monoxide Barium dioxide

This barium dioxide if heated to a higher temperature gives
up its excess of oxygen, and is reconverted into the original
barium monoxide.

2Ba02 = 2BaO + O2

Barium dioxide Barium monoxide

By thus alternately heating and cooling the barium com-
pounds, we can remove oxygen from the atmosphere and
then recover it in the pure state. As a matter of practice the
temperature at which the two reactions take place is not changed,
whereby a considerable saving in fuel is effected. The same
two actions can occur at constant temperature if the pressure
is varied. The air is first of all pumped under pressure
over the heated barium monoxide, which slowly absorbs the
oxygen. When the barium monoxide is charged with oxygen,
the pressure is then reduced. The effect of this is that the
oxygen which was absorbed under pressure is now given up,
and can be pumped into a suitable gas-holder, the barium
monoxide being regenerated, and once more made to take up
oxygen from the air under increased pressure. Here then
the temperature is kept constant, but the pressure is alter-
nately raised and lowered. Oxygen made in this way usually
comes into the market strongly compressed in steel cylinders.
Oxygen is prepared in the laboratory on the small scale by
heating some substance which gives up its oxygen readily.
The usual substance employed is potassium chlorate. When

OXYGEN 1^9

heated in a hard glass tube to a temperature somewhat below
a red heat, potassium chlorate fuses and gives up oxygen,
the equation for the complete action being —

2KCIO8 = 2KCI + 3O2
Potassium chlorate Potassium chloride

This same reaction may be carried out at a much lower
temperature by mixing a little manganese dioxide Mn02 with
the potassium chlorate. The manganese dioxide is not per-
manently changed at the temperature which is used to decom-
pose the chlorate, but it acts upon the chlorate in some way so
as to make the evolution of oxygen take place much more
rapidly and at a much lower temperature. A substance which
acts in this way — i.e, a substance which facilitates a chemical
action without itself undergoing any permanent change — is
called a catalytic agent. It should be noted that the oxygen
prepared from potassium chlorate is not so pure when
manganese dioxide is mixed with the chlorate, as when th6
chlorate is used alone.

Many other substances besides potassium chlorate give up
oxygen when heated. Thus potassium nitrate decomposes
as follows:—

2KNO8 * 2KNO2 + Oa

Potassium nitrate Potassium nitrite

Manganese dioxide, when heated to a high temperature, gives
up a third of its oxygen, according to the equation—

3Mn02 = Mn304 + O2

Red lead in a similar way decomposes as follows : —

2Pb804 = 6PbO + O2

Red lead Litharge

Oxygen prepared by any of these methods is a gas which
is somewhat heavier than air. Being colourless, it is invisible,
and it possesses neither smell nor taste. It cannot be com-
pressed to a liquid at the ordinary temperature, but when
sufficiently cooled it liquefies and produces a pale blue liquid.
It is, as we have seen, only very slightly soluble in water, and

I30 INORGANIC CHEMISTRY

in this connection it may be noted that solubility in water and
compressibility to the liquid form usually go hand-in-hand.
That is, gases which can be compressed to liquids at the
ordinary temperature are usually soluble in water, whilst
gases which are not compressible to liquids at the ordinary
temperature are usually but very slightly soluble in water.

All combustions which take place in air will also take place
in oxygen, the vigour of the chemical action and the brilliancy
of the combustion being greatly enhanced when pure oxygen
is used. Thus, a piece of sulphur or phosphorus will burn
with much greater brilliancy in pure oxygen than in air; and
a piece of iron wire, which, when heated to redness, soon
becomes cold in air, will produce showers of sparks when
introduced into a jar of oxygen, the combustion going on
until the metal is almost entirely converted into oxide. The
common test for oxygen is its action on a splinter of wood,
the flame of which has been extinguished so as only to leave
a tip of glowing carbon. When this is introduced into a
vessel of oxygen, the combustion of the carbon increases so
greatly in vigour that the splinter again bursts into flame.

Mixtures containing free oxygen in sufficient quantity will
always support combustion, and the vigour of the combustion
will be in general dependent on the amount of oxygen they
contain. Some gaseous compounds of oxygen also support
combustion, but the vigour of combustion in them is regulated,
not only by the proportion of oxygen contained in them, but
also by the ease with which the oxygen can part from the element
with which it is combined. Examples of this will be seen
when we consider oxides of nitrogen. It should be noted
that respiration cannot be supported except by free oxygen.
That is, a mixture of oxygen and nitrogen will support the
respiration of animals, whilst a compound of oxygen and
nitrogen will not, although it may easily support the ordinary
processes of high temperature combustion.

Ozone

A gas exists which differs in many ways from ordinary
oxygen, and yet contains nothing but the element oxygen. This
gas is known as ozone, and is said to be an allotropic modifica-
tion of oxygen.

OXYGEN

131

It is produced from ordinary oxygen under the influence of
electric excitation. When an electric machine is being worked,
a peculiar smell is perceptible in its neighbourhood. This
smell is due to ozone, which is produced from the oxygen of
the air in the neighbourhood of the electric spark. Ozone
may be produced in much larger quantity from oxygen by
passing the oxygen through a space which is subjected to
rapid alternation of electric charges, a tube devised for this
purpose being shown in the figure. The rapid change in the
electric charge is brought about by means of an induction coil,
the terminals of which become alternately positive and negative,
the rate of alternation being many hundred times a minute.

It is impossible by means of such a piece of apparatus to
convert the whole of any given amount of oxygen into ozone.

^=Br -

Fig. 27. — Preparation of Ozone.

Oxygen is passed through the ^lass tube A A', the metallic conductor S, which is
outside the tube, being connected with ong pole of an induction coil, and the conductor
C, which is inside the tube, being connected with the other. Under the influence of the
rapidly alternating electric charges a considerable proportion of the oxygen which
passes between B and C is transformed into ozone.

about one-fifth of the total oxygen being as much as can be
converted into ozone under ordinary conditions.

Ozone is a colourless gas which is about twice as soluble
in water as ordinary oxygen, and somewhat more readily
condensible to the liquid state. It differs from oxygen in
having a peculiar odour, and in the liquid produced from it
having a deep blue colour.

The chemical properties of ozone are also very different
from those of ordinary oxygen, although the difference is
rather one of degree than of kind. Ozone is a much
more powerful oxidising agent than ordinary oxygen. Thus,
at the ordinary temperature ozone will oxidise potassium
iodide with liberation of iodine, a reaction which is very
commonly used as a test for ozone, free iodine being

132 INORGANIC CHEMISTRY

readily visible from its colour, and still more so if starch
is added, blue iodide of starch being then produced. A
solution of a lead salt such as lead acetate is also oxidised
by ozone at the ordinary temperature, although it is quite
unaffected by oxygen. The lead salt is oxidised to lead
peroxide, Pb02, the change being evident by the production
of the brown colour of this compound. Indigo solution, which
retains its deep blue colour indefinitely in presence of ordin-
ary oxygen, is very soon oxidised by ozone, the blue colour
disappearing entirely. India-rubber is rapidly attacked by
ozone, and mercury loses its characteristic property of form-
ing a meniscus when it remains for any length of time in an
atmosphere containing ozone.

If we ask for an explanation of the difference in properties
that exists between ozone and ordinary oxygen, we find that
it is connected with a difference in the molecular weight of
the two gases. When a measured volume of ordinary oxygen
is partly converted into ozone, it is found that during con-
version a contraction in volume takes place, although the
weight of the gas remains the same. A given volume of ozone
must, therefore, be heavier than an equal bulk of oxygen — that
is, a greater weight of ozone can be contained at normal
temperature and pressure in 22.4 litres than is the case with
ordinary oxygen. The molecular weight of ozone must, there-
fore, be greater than that of ordinary oxygen.

It is not difficult to find by how much the molecular weight
of ozone is greater than that of oxygen. Ozone is entirely ab-
sorbed by oil of turpentine, whilst oxygen is quite unaffected by
this liquid. If we note the contraction which takes place
during partial conversion of a given volume of oxygen into
ozone, and then note the further contraction which occurs
when the ozone thus formed is absorbed by oil of turpentine,
we find that the second contraction is exactly double the first
contraction. If we call the original contraction one volume,
then the second contraction is equal to two volumes, and the
total contraction equal to three volumes. Now, the second
contraction is equal to the volume of ozone produced, and the
total contraction is the volume of the ordinary oxygen which
disappeared. As there was no loss of weight in the experi-
ment, it is evident then that three volumes of oxygen weigh as
much as two volumes of ozone under the same conditions of

OXYGEN 133

temperature and pressure. Put in another way, ozone is
volume for volume half as heavy again as oxygen, and there-
fore, according to Avogadro's principle, p. 114, its molecular
weight must be half as great again as that of ordinary oxygen
which is 32. The molecular weight of ozone is, therefore,
48, and its molecular formula is consequently O3.

When ozone is heated, it is reconverted into ordinary oxygen.
The equation for the conversion of oxygen into ozone, and for
the reconversion of ozone into oxygen, is as follows : —

3O2 ^ 2O

When ozone acts as an oxidising agent it is usually only the
extra oxygen which affects the oxidation, ordinary oxygen
being produced. Thus the equation which expresses the
action of ozone on potassium iodide is as follows : —

2KI + O3+H2O = 2KOH + O2 + I2

Potassium iodide Ozone Water Potassium hydroxide Oxygen Iodine.

Here the extra oxygen of the ozone disappears, and an equal
volume of ordinary oxygen is generated.

Ozone is said to exist in small quantity in the air.

CHAPTER XXI

WATEB

Wt must carefully distinguish between the chemical substance
water H2O, and the various kinds of water that occur in nature
— e,g, fresh water, salt water, or mineral water. No natural
water is the pure chemical compound. This is on account of
the solvent power of water, which, in the form of liquid, is
capable of dissolving at least traces of most substances with
which it comes into contact. Gaseous water and solid water —
t,e, water-vapour and ice — do not act as solvents in the same
way. If we wish, therefore, to obtain pure water from a
natural water containing dissolved material, the easiest
method is to convert it either into water-vapour or into ice,
and then liquefy these substances in vessels on which the
liquid water has no solvent action.

The purest liquid water that we find in nature is condensed
water-vapour in the form of rain. Even rain water, however,
contains substances in the state of solution, for rain in falling
through the air dissolves at the very least some of the gases
which constitute the atmosphere, and since the atmosphere is
never free from dust, the dust particles are also carried down
in some quantity with the rain. Rain water collected in towns
is, of course, far from pure, a great deal of soot and other
materials being carried down with it.

If the rain water falls on the ground and permeates the soil
it dissolves some of the constituents of the soil. Consequently,
we find that lakes, rivers, and springs (the water in all of
which is ultimately derived from rain water which has been
in contact with the earth's crust) all contain considerable
quantities of dissolved substances. The nature of the dis-
solved substances and the quantities of them in any given
water will, of course, depend upon the nature of the soil with
which the water has been in contact.

134

WATER 135

Waters from springs, which contain either very large
quantities of dissolved substances or else unusual substances .
which are easily recognised by the smell or taste, are called
mineral waters. The salt water which occurs in such
enormous quantities in the ocean contains, of course, common
salt in solution ; but besides common salt a great many other
substances are present. On the average sea water contains
between 3 and 4 per cent, of dissolved salts. The water of
salt lakes, such as the Dead Sea, contains even more dissolved
material, the Dead Sea holding over 20 per cent, of salts.
The ordinary tap water of town supplies is usually lake water
or river water, and does not in general contain much dissolved
solid, the quantity not being more than about one-hundredth
of a per cent.

Two of the commonest substances in ordinary water supplies
are calcium hydrogen carbonate CaH 2(003)2, and calcium sul-
phate CaS04. If a water contains much of these substances
it is said to be a liard water; if it contains little of these
substances it is said to be a soft water. Hard waters are
unsuitable for washing purposes on account of the amount of
soap they render useless before a lather is formed. Soap is a
mixture of the sodium salts of various acids of which the
calcium salts are insoluble. When, therefore, we bring the
soap into contact with the water containing dissolved calcium
salts, double decomposition takes place, sodium hydrogen
carbonate and sodium sulphate being produced on the one
hand, and the insoluble calcium salts of fatty acids on the
other. The insoluble calcium salts fall out of the solution,
whilst the sodium sulphate and sodium hydrogen carbonate
remain dissolved. These, however, are quite valueless as
cleansing agents, and so more soap must be added until all
the calcium salts are precipitated. The degree of hardness of
water (usually stated in parts of calcium carbonate per 100,000)
is generally ascertained by finding how much of a standard
soap solution the water will use up before it produces a
permanent lather — Clarke's soap test.

The hardness produced by calcium hydrogen carbonate
differs from the hardness produced by calcium sulphate.
The former is called temporary liardness, the latter per-
manent liardness. These names are applied on account
of the hardness caused by the calcium hydrogen carbonate

136 INORGANIC CHEMISTRY

disappearing on boiling, while the hardness caused by the
calcium sulphate persists after the water has been boiled.
The action which results in the disappearance of the temporary
hardness on boiling is the following : —

f
CaH2(C03)2 - CaCOs + HgO + CO2

The carbon dioxide escapes as gas, and the normal calcium
carbonate falls out as solid. If the water considered has
much calcium hydrogen carbonate in solution, it deposits a
fur or crust on the interior of the vessel in which the hard
water has been boiled. This crust is usually very hard, and is
a poor conductor of heat, so that much more fuel is required
to boil water in a vessel which has been thus encrusted than
to boil water in a clean metal vessel. This renders such
water unsuitable for steam-raising purposes.

If we wish to estimate both temporary and permanent hard-
ness by the soap test, we first of all take a sample of the water
and estimate its total hardness. Then we boil a sample and
estimate the hardness which remains after boiling. The
difference between the total and permanent hardness gives
the hardness which has disappeared — that is, the temporary
hardness.

Rain water, since it contains little or no dissolved solid, is
very frequently used for washing purposes, where the other
natural waters are hard. Hard water, however, may be
softened in various ways. For washing purposes the common
substance to employ is washing soda or sodium carbonate.
When a little of this is added to the hard water, the following
reactions take place : —

NagCOa + CaH2(C03)2 = 2NaHC03 + CaCOs
NagCOg + CaSO^ = Na2S04 + CaCO g

Calcium carbonate is precipitated and nothing but sodium
salts are left in the water, and these in no way impair the
cleansing power of soap. Another substance which may be
used for softening water is calcium hydroxide. When this is
added to a solution of ^alcium hydrogen carbonate^ normal

WATER 137

insoluble calcium carbonate is produced according to the
equation —

Ca(OH)2 + CaH2(C03)2 = 2CaC03 + 2H2O

and thus the water loses its temporary hardness.

The pure chemical substance H2O is colourless when
viewed in a thin layer, but is seen to be greenish blue
when a thick layer is traversed by a beam of light. It is
without odour, but possesses a somewhat mawkish taste.
The purest water we employ in the laboratory is distilled
water — i.e, condensed steam, — and for all practical purposes
this distilled water may be treated as pure water. It is a
standard substance for many purposes, because it can be
easily prepared in a state nearly approaching purity, and can
be obtained in unlimited amount. Thus, the zero point on the
centigrade thermometer is fixed at the temperature at which
pure water freezes, and similarly 100** on the same thermometer
is fixed as the temperature at which pure water boils under a
pressure of 760 mm. Pure water is also taken as the
standard of specific gravity, the specific gravities of all other
substances being practically referred to water under specified
conditions as unity.

Water is a substance distinguished for its solvent power
and it is by far the commonest solvent we employ. Nearly
all the substances we use in the chemical laboratory are
dissolved in water if they are at all soluble. The reason for
this has already been given on p. 9. Water is distinguished also
by its power of ionising salts, acids, and bases, dissolved in
it. No other common solvent possesses this ionising power
to anything like the same degree.

Water has very different solvent properties according to the
substances with which we bring it in contact, and also accord-
ing to the temperature at which it acts on these substances
(p. 15). It should be noted that the substances which we
usually speak of as being insoluble are in reality slightly
soluble in water. The slight solubility may not be of im-
portance in the chemical laboratory, but it is sometimes of
great importance in the phenomena of nature. Thus, as we
have seen, water will take up from the atmosphere about 0.8 per
cent, of its own volume of oxygen, or if we measure by weight
instead of by volume, about 0.00 1 per cent, of its own weight

138 INORGANIC CHEMISTRY

of oxygen. This quantity is very small, but it is on this small
amount of oxygen that fish and other animals which inhabit
the water have to depend for their respiration. Rocks, too,
which we should cdl insoluble in water, are in reality
gradually dissolved away by water when the water acts over
great lengths of time. Glass is a substance which we commonly
employ to contain water and aqueous liquids, and we usually
state it to be insoluble in water. It is easy to show, how-
ever, that water when boiled in glass vessels attacks them
and dissolves up recognisable quantities. The so-called in-
soluble precipitates are by no means insoluble in water;
thus, the following substances, which are generally treated
as insoluble substances, dissolve in water to the extents given
below : —

Milligrams dissolved by
Substance i litre of water at 18^.

Silver chloride .
Mercurous chloride
Barium sulphate
Lead sulphate .
Calcium carbonate

1.7

2.6
46

Many substances when they separate out from solution in
water, separate out with what is called water of crystallisation.
This water of crystallisation is combined with the substance
in some way which we do not altogether understand, and is
in no sense liquid water. Ordinary washing soda, for example,
contains water of crystallisation, and if we wish to represent
the amount of the water of crystallisation, we can do so
by means of the formula Na2C03,ioH20. The sodium
carbonate and the water of crystallisation, are contained in
washing soda in perfectly definite proportions, and so it is
with other substances of a similar nature. Substances which
contain water of crystallisation are called liydrates, and we
can indicate the number of formula-weights of water combined
with one formula-weight of substance by means of the Greek
numerals. Thus, washing soda NagCOsjioHgO is called a
decahydrate of sodium carbonate ; blue vitriol^ CuS04,5H20
is called a pentahydrate of copper sulphate ; Glauber's salt
Na2S04,7H20 is a heptahydrate of sodium sulphate, and

WATER 139

so on. These hydrates part with their water when heated.
Very frequently the water comes off at ioo°C., but occasionally
the water may remain combined with the substance at a
much higher temperature than this. Thus, if we heat blue
vitriol CuS04,sH20, four of the five formula-weights of water
are easily driven off by heat, while the fifth formula-weight of
water is not driven off until the temperature is raised to a
much higher point. Hydrate-forming substances when without
water of crystallisation are said to be anhydrotui.

Some hydrates lose their water of crystallisation when exposed
to the air at the ordinary temperature. This may be observed
with a clean crystal of washing soda. The surface of the
crystal is at first bright and uniform, but soon becomes
covered on exposure to the air with a white powder,
especially at the angles of the crystal. This white powder is a
hydrate of sodium carbonate, which contains less water than
the washing soda, and is formed from the washing soda by loss
of water to the air. Substances which behave in this way are
called efflorescent, and the crust that appears on the surface
is called an efflorescence.

Other substances behave in exactly the opposite way. An an-
hydrous substance or a lower hydrate may absorb moisture from
the air to form a higher hydrate. Such hygroscopic substances
may be used as drying agents.

The crystals of calcium chloride hexahydrate, which have
the formula CaCl2,6H2 0, continue to absorb water from
the air, and finally pass into solution in the water which
they have absorbed. Substances of this kind are said to be
deliquescent. The tendency of anhydrous calcium chloride
to absorb water is so great that it is very frequently employed
to dry gases and liquids, which do not themselves act upon
the calcium chloride. The calcium chloride absorbs the water
in these moist gases or liquids, and becomes first of all a
hydrate of calcium chloride, and then, if sufficient water
is present, a concentrated solution of calcium chloride. It
should be noted that all deliquescent substances are very
soluble in water. A substance which is insoluble in water or
only moderately soluble in water can never be deliquescent in
ordinary air.

The hydrates have very often different colours from the
anhydrous substances, and use is made of this fact in the

I40 INORGANIC CHEMISTRY

preparation of sympathetic inks, A dilute solution of cobalt
chloride CoCU bas a pale pink colour, and wben used as ink
dries up on the paper to a practically colourless hydrate.
When this hydrate is heated, however, by holding the paper
before the fire, the water of crystallisation is driven off and
the anhydrous chloride which remains is plainly evident from
its deep blue colour. When the paper is allowed to cool, the
blue anhydrous salt again absorbs moisture from the air and
the writing disappears.

The chief chemical properties of water have already been
discussed. Thus it frequently unites with basic oxides to form
hydroxides —

CaO + H2O = 0(OH)2

and with acidic oxides to form acids —

SO3 + H2O = H2SO4

It is decomposed by many metals : by some at the ordinary
temperature, thus —

2Na + 2H2O = H2 + 2NaOH;
and by others only at a high temperature, thus —
Fe + H2O = H2 + FeO

hydrogen being formed on the one hand, and a metallic
hydroxide or oxide on the other. Not only is it decomposed
by metals at a high temperature; it is also decomposed by
some non-metallic elements. For example, we have seen
that when brought into contact with carbon at a red heat, it is
decomposed with formation of hydrogen and carbon monoxide,
according to the equation —

C + H2O = H2 + CO

We have seen, too, that when some salts, bases, and acids are
dissolved in it (for example sodium sulphate, sodium hydroxide,
and sulphuric acid) it can be easily decomposed by means of
an electric current into hydrogen and oxygen (Chapter XIV.).
In connection with this, it possesses in a very high degree the
power of ionising salts, acids, and bases dissolved in it, splitting
them up to a greater or less extent into their positive and
negative radicals each with its appropriate electric charge.

WATER 141

Not only has it this power of converting salts into their ions ;
it has also the power of partially decomposing many salts into
the acid and base from which they are derived by neutralisation.
It has already been pointed out (p» 65), that only normal salts
derived from strong bases and strong adds, are neutral to in-
dicators when they are dissolved in water. A salt derived from
a strong base with a weak acid has an alkaline reaction. ^Thus,
sodium carbonate, which is derived from the strong base, sodium
hydroxide NaOH, and the weak acid, carbonic acid HgCOs, ^s
alkaline to litmus and many other indicators. The same is
observed with a solution of sodium sulphide, derived from
sodium hydroxide and the weak acid sulphuretted hydrogen.
On the other hand, salts such as copper sulphate, ferric chloride,
and aluminium sulphate (or alum), which are derived from
strong acids and the weak bases Cu(0H)2, Fe(0H)3, A1(0H)3,
give aqueous solutions which have always an acid reaction
to ordinary indicators. The reason why these solutions are
alkaline or acid, is that they actually contain free alkali and
free acid respectively, produced by the action of the solvent
water on the normal salt. This action of water on a normal
salt, with production of free acid or free base, is usually termed
hydrolysis.

In the case of sodium carbonate, the action of the water may
be represented by the reversible equation —

Na^COg + H2O ^ NaHCOs + NaOH

It must be understood that in most cases of this kind the
hydrolysis only proceeds to a small extent. In a weak solution
of sodium carbonate only about i per cent, of the total
dissolved salt is thus split up into sodium hydroxide and
the bicarbonate. The sodium hydroxide is strongly alkaline
in its action, and sodium hydrogen carbonate practically
neutral, so that the solution on the whole has an alkaline
reaction.

The hydrolysing action in the case of ferric chloride may be
represented as follows : —

FeCls + 3H2O = Fe(0H)3 + 3HCI

This action also goes on to a very small extent, but the hydro-
chloric acid produced is a strong acid, and the ferric hydroxide

142 INORGANIC CHEMISTRY

produced is a very weak base, so that the solution on the
whole has a distinctly acid reaction. Ferric hydroxide is what
we ordinarily call an insoluble substance. In the above case,
however, it does not fall out of the solution, but remains sus-
pended in the solution in such a state that it cannot be
detected by the eye, although its presence in the undissolved
state may be rendered evident by proper optical means.
A substance in this condition is said to be in a state of psetuto-
solution.

CHAPTER XXII-

CABBON

At.l living substances contain the element carbon, which is
also a constituent of a great number of minerals, chiefly
carbonates.

The element carbon itself is found in nature in a more or
less pure condition. It exists in two crystalline modifications :
first, the comparatively rare and precious diamond ; second, the
much more common and less valuable grapMte. These two
varieties of carbon when pure contain nothing but the element
carbon, yet they have absolutely different physical properties,
although their chemical properties are practically speaking
identical. Diamond is the hardest substance known j graphite,
which is commonly called plumbago or black leady is so soft
that it is used in the manufacture of writing pencils — Le, it is
so soft as to be abraded by paper and leave a track of black
particles behind. Diamond is colourless and transparent;
graphite is black and opaque.

When either of these substances is heated in oxygen it burns
to form carbon dioxide, and a given weight of either substance
will, if pure, yield a quantity of carbon dioxide in the proportion
expressed by the equation —

C + O2 = CO2

This experiment shows that the two substances, although
differing so greatly in their physical properties, are chemically
identical.

There are many varieties of artificially prepared carbon, all
more or less impure. Lamp-black is one of these, and is
simply condensed smoke. Certain substances, like oil of
turpentine and acetylene, burn with a very smoky flame wheri
an insufficient supply of air is used in the combustion. The
smoke produced consists of small particles of unburnt carbon,

143

144 INORGANIC CHEMISTRY

and may be made to deposit on the walls of a chamber or on
sheets, which, when scraped, give carbon in the form of loose
powder known as lamp-black. Lamp-black is chiefly used in
the production of printers' ink, and as an ingredient of certain
pigments.

Another form of carbon which is extensively used is charcoal.
This charcoal is by no means so pure a form of carbon as
lamp-black, but it is produced in much greater quantities. It
i^ made by heating wood to such a temperature that most of
the organic substances which form the wood are decomposed
and driven off as gases, a residue of impure carbon remaining
behind.

The production of charcoal was formerly carried out entirely
by piling up billets of wood, covering the pile with turf, and
setting fire to the wood at the bottom of the pile, the air supply
being carefully regulated. Part of the wood burned in the
defective supply of air, and the whole of the mass was heated
by this partial combustion of the wood to such a temperature
that charcoal was left behind. The kind of charcoal obtained
depended on the temperature to which the wood was heated.
All charcoal is more or less porous and light, as it retains
the original form of the wood, although most of the material
of the wood has been removed in the form of gas. If the
charcoal burning is conducted at a high temperature, the char-
coal is comparatively dense and contains very little of other
substances than carbon and the incombustible salts forming
the ash of the wood; whilst^ if the temperature is low, the
charcoal still contains considerable proportion of carbon com-
pounds as well as carbon itself. Nowadays a great deal of
wood charcoal is prepared by heating the wood in retorts, air
being excluded from the process altogether. The wood is
decomposed by the rise of temperature as before, but instead
of the gases and vapours derived from the wood being allowed
to burn, they are carefully collected and a variety of useful
products are derived from them.

On account of its porous nature and the very large internal
surface which it presents, charcoal has the power of condensing
large amounts of gases. Thus, if a piece of dry charcoal is
introduced into a cylinder of ammonia contained over mercury,
the ammonia will be absorbed by the charcoal and the mercury
will rise in the cylinder to take its place. Similarly, charcoal.

CARBON

especially bone-charcoal prepared by heating bones, will absorb
colouring matters from solutions, so that it finds a consider-
able application as a decolorising and deodorising agent. The
chief use of wood charcoal is as a smokeless fuel and as an
ingredient of gunpowder.

Goal is a natural variety of impure carbon which occurs in
enormous quantities. Coal is formed by the decomposition
of vegetable matter in a special way. When wood decays in
absence of air it loses a portion of its carbon and proportionately
more of its hydrogen and other gaseous elements, so that the
residue contains more carbon than the original wood. The
following table gives roughly the percentage of carbon in
dry wood, and of some kinds of vegetable matter which have
undergone partial decay: —

Carbon
Dry wood . . . • 5© per cent.

Peat .
Brown coal .
Cannel coal .
Newcastle coal
Anthracite .

60
70

85
90

When this kind of decay has gone on for a very long time
coal is the result, and coal contains amounts of carbon
varying from 80 to nearly 100 per cent.

Coal finds its chief use as fuel and in the production of
coal-gas, but it is also used as a chemical agent in many
reducing processes of which we have already had examples.

A form of carbon which is much used by electricians is what
is known as gas carbon. This is a hard substance which is
formed on the walls of retorts in which coal is heated for the
manufacture of coal gas, and is a good conductor of electricity.
To form the carbons for electric lighting this gas carbon is
ground to powder and then moulded with tar into the proper
form and strongly heated.

Coke is another form of carbon, also found in gas retorts.
It is formed by heating coal to a very high temperature, the
volatile gases which come off from the retorts being partially
condensed to coal tar and partially used for illuminating
purposes as coal-gas. The coke, which contains a portion
of the original carbon and a large quantity of ash, remains in

146 INORGANIC CHEMISTRY

the retort, and is used chiefly as fuel. It burns with difficulty,
but produces a clear, smokeless fire.

It has already been noted that carbon burns without flame,
because it cannot be converted into gas at the temperature
which can be generated by its own combustion. Indeed,
carbon cannot even be melted by the temperature of its own
combustion, and can only be very partially softened by the
intense heat of the electric arc.

Oxides of Carbon

Carbon forms two oxides, carbon monoxide CO and carbon
dioxide CO 2. Although carbon monoxide may probably be
formed in small quantity in the direct combustion of carbon
when the supply of oxygen is defective, there is no doubt that
the greater part of the carbon which combines with oxygen
is directly converted into carbon dioxide.

Carbon dioxide, CO 2, is not only evolved in the processes
of combustion, respiration, and fermentation, but is also
produced in large quantities by volcanoes, so that it is a
constant constituent of the atmosphere. It is a gas which is
already fully supplied with oxygen, and will not burn in the air.
The oxygen which it contains, however, is held in firm union
by the carbon, so that it does not act readily as an oxidising
agent or a supporter of combustion. If a lighted taper or
other substance is plunged into a jar filled with the gas it is
at once extinguished. Carbon dioxide, as we can calculate
from its formula-weight, is considerably heavier than air, and
is generally collected by downward displacement. It can be
poured from one vessel to another like a liquid. A simple test
for its presence is the turbidity which it produces in coming
into contact with lime-water or baryta water.

The ordinary mode of producing carbon dioxide in the
laboratory is to treat calcium carbonate, usually in the form
of marble, with hydrochloric acid. Decomposition occurs
according to the equation —

CaCOs + 2HCI = CaClj + HgO + CO2

and the carbon dioxide, being the only gaseous product, is
obtained nearly pure. A machine which is frequently used

CARBON

147

in the laboratory for the production of carbon dioxide in this
way is that described in fig. 28.

Carbon dioxide is a gas which can be compressed to a liquid
if sufficient pressure is employed,
and liquid carbon dioxide is sup-
plied in steel cylinders under
a pressure of about fifty atmos-
pheres. When prepared on the
large scale, it is either made from
some form of calcium carbonate,
by treatment with acid or by
heat alone, or else it is obtained
as a by-product in the process
of fermentation.

Corresponding to the compara-
tively easy compressibility to a
liquid, we have moderate solubility
of carbon dioxide in water. At
the ordinary temperature water
can dissolve about its own vol-
ume of carbon dioxide. Some
of this carbon dioxide remains
unchanged in the water, but some
of it unites with the water and
becomes carbonic acid, according
to the equation —

H2O -♦- CO2 = H2CO3

Carbonic anhydride Carbonic acid

Effervescing beverages are aqueous
liquids charged with carbon dioxide
under pressure. They usually con-
tain about five times as much car-
bon dioxide as the liquid would be
capable of dissolving at one atmos-
phere pressure. In order that the
liquid may retain this quantity of

carbon dioxide, the pressure within the bottle must be equal
to about five atmospheres. When the bottle is uncorked the
pressure over the liquid is at once reduced to one atmosphere.
Under these conditions the liquid can only retain in the dis-

Fig. 28. — Kipp Machine.

This piece of apparatus is usi*<i
to obtain an automatically regulated
supply of a gas, e^. carbon di-
oxide, produced by tne inter-action
of a liquid and a solid, e,g. hydro-
chloric acid and marble. There is
no connection between A and C
except through B. When the tap
•S" is opened, the acid falls in A and
rises in B until it comes in contact
with the marble in C. If the chemi-
cal action thus induced supplies
more gas than can escape through
.S", the acid is forced downwards
away from the marble, so that less
gas is generated. In this way the
apparatus is self-regulating.

148 INORGANIC CHEMISTRY

solved state about one-fifth of the amount it originally contained,
and the remaining four-fifths escape from the liquid in the form
of bubbles. The water of some natural springs is slightly
charged with carbon dioxide, which has been dissolved under
considerable pressure at a distance below the earth's surface.

Rain in falling through the atmosphere absorbs some
carbon dioxide, which, when dissolved by the water, is jjartially
converted into carbonic acid. The carbonic acid thus formed
enables the water to dissolve many substances which it would
not otherwise attack. In particular, it enables the water to
dissolve the insoluble normal carbonates. Thus calcium
carbonate, which is practically insoluble in pure water, dis-
solves to a much more considerable extent in water contain-
ing carbonic acid. In this solution, however, it does not
exist as the original normal calcium carbonate, but is con-
verted by the carbonic acid into the soluble acid carbonate,
in accordance with the equation —

CaCOg + HgCOs = CaH2(C03)2

this acid carbonate conferring temporary hardness on the
water (p. 135). The carbonic acid of the air therefore plays
a great part in the disintegration of rocks, particularly those
consisting of carbonates.

Carbon dioxide is not in the ordinary sense a poisonous
gas, but when mixed with air in the proportion of one to
ten it may cause death by suffocation, chiefly on account of
its preventing the carbon dioxide brought by the blood to
the lungs from escaping freely into the air, and thus the
corresponding volume of fresh air from taking its place.
The choke-damp in mines which produces suffocation is chiefly
due to the presence of carbon dioxide.

Carbon monoxide, CO, differs greatly in its properties, both
physical and chemical, from carbon dioxide. It is formed,
as we have seen, by the combustion of carbon in a defective
supply of oxygen, chiefly by the action of carbon dioxide on
excess of heated carbon, the equation being —

C + CO2 = 2CO

As it is not fully saturated with oxygen it bums in air with
a blue flame to produce carbon dioxide, and can act as

CARBON

149

a reducing agent. The oxygen which it contains is firmly
held, so that it is not itself a supporter of combustion. It is
a gas which is slightly lighter than air, as its formula-weight
indicates, and cannot therefore be collected by downward dis-
placement. Unlike carbon dioxide, it is not easily liquefied,
no pressure, however great, being capable of reducing it to the
liquid state at the ordinary temperature. Corresponding to
this, it is only very slightly soluble in water.

When breathed, even in small quantity, it acts as a powerful
poison, combining with the haemoglobin of the blood, and
preventing it from absorbing oxygen in the lungs. It is the

Fig. 29. — Preparation of ** Water-gas" and Carbon Monoxide.

Steam is generated hy boiling water in the flask /^, and is passed over pieces of
charcoal contained in the iron tube 7^, which is heated to bright redness in the charcoal
furnace C. The gas, which is collected over water, consists of carbon monoxide and
hydrogen. If carbon dioxide is delivered into the tube instead of steam, the gas which
issues is nearly pure carbon monoxide.

poisonous gas contained in the fumes from a bright charcoal
fire, and it is also to its presence that the poisonous character
of coal-gas is due.

When air is passed over white-hot carbon, the oxygen is all
consumed by the carbon and converted into carbon monoxide,
the nitrogen of the air not being attacked. This mixture of
carbon monoxide and nitrogen is called producer gas^ or
generator gas, and is sometimes used as a reducing agent
and as a gaseous fuel.

Another kind of gaseous fuel may be obtained by passing
steam over red-hot carbon. The carbon combines with the

I50 INORGANIC CHEMISTRY

oxygen of the water to form carbon monoxide, and the
hydrogen of the water is liberated.

H«0 + C = CO + H

A mixture of gases thus obtained is usually called water gas,
and contains nitrogen, carbon monoxide, and hydrogen. The
last two gases are both combustible, so that water gas can
either be used as a gaseous fuel, as a reducing agent, or, when
enriched by naphtha or other suitable substance, as an
illuminating gas.

Mond gas is produced by the joint action of steam and air
on heated coal-slack. It contains about half its volume of
nitrogen, the other half being a mixture of hydrogen, carbon
monoxide, and carbon dioxide. It is mostly used as a fuel for
gas-engines.

Hydrocarbons

The compounds which contain only carbon and hydrogen
are called hydrocarbons. Whilst carbon only forms two com-
pounds with oxygen, it forms hundreds of compounds with
hydrogen. We shall only consider here a few of the simplest
gaseous hydrocarbons.

Methane or marsh gas, CH4, contains the largest proportion
of hydrogen of any hydrocarbon. When vegetable matter
decays beneath the water of a marsh, methane is produced, and
rises in bubbles to the surface, whence the name marsh gas. It
cannot be formed by the direct union of carbon and hydrogen ;
indeed, these two elements can only be made to unite under
exceptional circumstances. It can be made, however, indirectly,
and is a product of the decomposition by heat of a great
many organic substances. It is one of the chief constituents
of coal-gas, made by heating coal to a high temperature
in closed retorts. An average sample of coal-gas contains
about one-third of its volume of marsh gas. The gas which issues
from the earth^s crust in enormous quantities in the neighbour-
hood of oil wells consists chiefly of marsh gas ; and it is also
found enclosed in fissures in coal-seams, so that when these
enclosures are broken into by miners, the marsh-gas escapes
into the mine, and there forms a combustible mixture with the
air in the mine. It is. known for this reason as fire-damp.

CARBON 151

Coal-mine explosions are mostly due to the ignition of this
explosive mixture.

Like all hydrocarbons it burns readily in air, the complete
combustion being expressed by the equation —

CH4 + 2O2 = CO2 + 2H2O

I vol. 2 vols.

If marsh gas and oxygen are taken in these proportions by
volume, the mixture is violently explosive, especially when in
large quantities.

Marsh gas bums in air with a flame which is only slightly
luminous. It is not poisonous, is incompressible to a liquid
at the ordinary temperature, and corresponding to this, is only
slightly soluble in water.

Ethylene, C2H4, is formed along with marsh gas during the
dry distillation of coal, although not in anything like the
same proportion. It burns with a bnghtly luminous flame,
and is one of the illuminating agents present in coal-gas.

It may be prepared in the pure state by the action of con-
centrated sulphuric acid on alcohol. If these substances are
warmed together in a flask, a brisk reaction takes place,
whereby alcohol loses water, according to the equation —

CaHeO = C2H4 + H2O

Alcohol Ethylene Water

ethylene gas being produced. At the same time, some of the
sulphuric acid acts as an oxidising agent on the alcohol and
destroys it, the sulphuric acid being, under these circumstances,
reduced to sulphur dioxide. The ethylene which comes off,
therefore, is mixed with sulphur dioxide as well as traces of
alcohol vapour. It may be freed from these impurities by
passing the gas first through strong caustic soda solution, and
then through water. The sulphur dioxide is converted by the
caustic soda into sodium sulphite, and the alcohol is retained
by the water. This impurity may be avoided by the use of
syrupy phosphoric acid instead of sulphuric acid.

The ethylene produced in this way is, like methane, a colour-
less combustible gas, the equation for its combustion being —

C2H4 + 3O2 = 2CO2 + 2H2O

I vol. 3 vols.

It has a faint pleasant odour, can be compressed to liquid at

IS2 INORGANIC CHEMISTRY

the ordinary temperature, and is much more soluble in water
than marsh gas.

Ethylene is what chemists call an nnsaturated substance,
a term which is usually applied to such carbon compounds
as can combine directly with chlorine or bromine. If ethylene
gas is led through a tube containing bromine, the two sub-
stances unite, according to the equation —

C2H4 + Br2 = C2H4Br2

Ethylene Ethylene dibromide

with formation of ethylene dibromide^ which is a heavy,
colourless, fragrant oil. Similarly, if ethylene and chlorine gas
are brought together in equal volumes, they unite, with pro-
duction of ethylene dichloridey the equation being —

C2H4 + CI2 = C2H4CI2
1 vol. I vol. Ethylene dichloride

This reaction was observed long ago by certain chemists in
Holland, and ethylene dichloride was for that reason called
oil of Dutch chemists, and the ethylene which produced it
is still frequently known under the name of olefiant gas — i.e,
oil-producing gas.

Acetylene, GI2H2, is a hydrocarbon which has of late assumed
great practical importance. It is formed when coal-gas burns
in a defective supply of air, and is probably produced in the
interior of all hydrocarbon flames, being afterwards burned
to carbon dioxide and water, when it reaches the more plentiful
oxygen supply in the external flame. It has long been known
for the extremely luminous character of the flame which it
produces when burnt in air, but it is only lately that a process
has been found for its manufacture in large quantity.

When quicklime is raised to the intense heat of an electric
furnace along with carbon, usually in the form of coal-slack,
the following action takes place : —

CaO + 3C * CaC2 + CO

Calcium oxide Carbon Calcium carbide Carbon monoxide

the products being carbon monoxide, which comes off as gas,
and calcium carbide, which remains in the furnace, in the
form of a hard grey crystalline mass. When treated with

CARBON 153

water, calcium carbide is at once decomposed with production
of calcium hydroxide and acetylene.

CaCg + 2H2O = Ca(OH)2 + C2H2

It is therefore an easy matter to prepare acetylene gas from
calcium carbide, and all acetylene generators are merely ap-
pliances to produce an automatically r^ulated supply of
acetylene gas by the interaction of calcium carbide and
water.

The equation for the combustion of acetylene is as follows : —

2C2H2 + 5O2 = 4CO2 + 2H2O

2 vols. 5 vols.

If acetylene and oxygen are mixed together in these propor-
tions, and a light is applied, the mixture explodes with great
violence, usually shattering the vessel in which it is contained.

Acetylene, like ethylene, is produced in the distillation of
coal, although only in small quantity, and helps to give coal
gas its illuminating power. Like ethylene, it is an unsaturated
substance, combining directly with chlorine and bromine to
form oily liquids.

When pure, acetylene has a smell which is not at all un-
pleasant, and closely resembles that of ethylene. Impure
acetylene, derived from calcium carbide, on the other hand,
has an obnoxious smell, and very poisonous properties. Both
the unpleasant smell and the poisonous character of this com-
mercial acetylene may be got rid of by suitable methods of
purification.

Acetylene can easily be distinguished from the preceding
hydrocarbons by its action on a solution of silver nitrate to
which ammonia has been added. The pure acetylene when
bubbled through this solution produces a white precipitate,
while ethylene and marsh gas do not affect the solution at
all. The precipitate is an insoluble substance called silver
acetylide. Crude commercial acetylene, instead of producing
a white precipitate, produces a black precipitate owing to the
presence of impurities, such as sulphuretted hydrogen and
phosphine.

CHAPTER XXIII

NITROGEN

Nitrogen, as we have seen, forms the bulk of the atmosphere.
It is an extremely inactive substance, and as an element only
enters into few chemical actions, although many of its com-
pounds are extremely active. It has already been said that
what was formerly known as atmospheric nitrogen is not
pure nitrogen, but a mixture containing about i per cent
of argon together with much smaller quantities of other inert
gases.

Pure nitrogen may be prepared from compounds of nitrogen.
Thus, if any of the oxides of nitrogen is passed over heated
copper, the copper decomposes these oxides, retaining the
oxygen and liberating nitrogen.

2Nq + 2Cu = 2CuO + N2

Nitric oxide

N2O + Cu = CuO + N2

Nitrous oxide

This method may be used in the laboratory, but if only a
small quantity of nitrogen is required, the following may be
substituted; Ammonium nitrite when heated splits up accord-
ing to the following equation : —

NH4NO2 = 2H2O + N2

Ammonium nitrite

It is not necessary to prepare ammonium nitrite, which is not
a common salt, specially for this purpose, since the action
takes place equally well if we heat a solution containing
potassium nitrite and an ammonium salt — e.g, ammonium
sulphate. Potassium sulphate and ammonium nitrite will be
formed by double decomposition, and the ammonium nitrite

154

NITROGEN 155

^ill be at once decomposed on heating, so that the whole
action is represented by the equation —

(NH4)2S04 + 2KNO, = K2SO4 + 4H2O + 2N2

Ammonium Potassium

sulphate nitrite

The nitrogen may be removed from the atmospheric mixture
of nitrogen and argon by either of the following methods : — If
excess of oxygen is added to the mixture, and a series of
powerful electric sparks passed through the moist gases, the
nitrogen, oxygen, and water react so as to produce nitric acid,
which can be absorbed by means of alkali. The process is
a comparatively slow one, but if carried on for a sufficiently
long time, all the nitrogen may be made to disappear, a
mixture of argon and oxygen remaining behind. The oxygen
may easily be removed by means of metallic copper, and so
aigon may be obtained. The other method is to pass the
mixture of nitrogen and argon over heated metallic magnesium.
Magnesium under these circumstances combines slowly with
the nitrogen to produce a substance known as magnesium
nitride.

3Mg + N2 = NgMgs

Magnesium Magnesium nitride

The nitrogen can thus be all absorbed, the less active argon
being unattacked. The only process in which atmospheric
nitrogen is used on the large scale is in the production of
cyanides such as barium cyanide BaC2N2.

BaCg + N2 = Ba(CN)2

Barium carbide Nitrogen Barium cyanide

Nitrogen is one of the least soluble of gases, and in connec-
tion with this, it is very difficult to condense to the liquid
state, the temperature having to fall below - 140" before the
condensation can be effected by pressure.

Ammonia, NH3

Nitrogen forms several compounds with hydrogen, of which
ammonia is by far the most important. The elements cannot

156 INORGANIC CHEMISTRY

be made to unite directly in any considerable quantity, but
ammonia, when once formed, is a stable substance.

In the distillation of coal to produce coal-gas, much of the
nitrogen contained in the coal comes off with hydrogen in the
form of ammonia, which is for the most part converted at the
gas-works into ammonium sulphate. From this, or any other
ammonium salt, ammonia may readily be prepared by mixing
it with lime and heating, the action which takes place being as
follows : —

CaO + (NH4).,S04 = CaS04 + 2NH3 -h H2O

The ammonia may be dried by passing over a layer of quick-
lime, which absorbs any water vapour it may contain.

Ammonia is a gas which is easily detected by its characteristic
smell. It is considerably lighter than air, and may be collected
by upward displacement. It is excessively soluble in water,
and can be liquefied by pressure at the ordinary temperature
of the atmosphere. The solution in water contains a large
proportion of ammonium hydroxide, formed according to the
equation—

NH3 -h H2O = NH4OH

All the ammonia may be expelled from solution in water by
continued boiling.

Ammonia burns readily in an atmosphere of oxygen —

4NH3 + 3O2 = 2N2 + 6H2O

but not in air unless heat is supplied to keep the action going
(p. 10).

When a succession of electric sparks is passed through
ammonia gas it gradually is decomposed into its elements,
the equation being —

2NH3 = N2 + 3H2

2 vols. I vol. 3 vols.

Ammonia is not in itself a base, and only becomes one on
combination with water. It may, however, be looked on as a
kind of anhydrous base, since it has the power of neutralising
acids with formation of ammonium salt but without forma-

NITROGEN 157

tion of water. This may be seen by comparing the two
equations —

NH4OH + HQ = NH^a + HaO
NH3 + HCl = NH4CI

Oxides of Nitrogen

Nitrogen, although it does not combine with oxygen directly,
can be obtained in combination with oxygen in the form of
five different oxides, which are all produced from nitric acid
or the nitrates.

Nitrogen pentozide, N2O5

Nitrogen pentoxide may be formed from nitric acid by simple
abstraction of water. One of the most powerful dehydrating
agents that we know is phosphorus pentoxide P2O5. When
warmed with nitric acid, this substance abstracts from it
the elements of water, according to the equation —

2HN08

+ P*Os

= 2HPO8 +

N,Os

Nitric acid

Phosphorus

Metaphosphoric

Nitrogen

pentoxide

acid

pentoxide

and nitrogen pentoxide passes off as gas. By cooling, it
can be condensed to a white solid which unites greedily with
water to reproduce nitric acid.

N2O5 + H2O = 2HNOs

Nitric anhydride Water Nitric acid

Nitrogen pentoxide is therefore the anhydride of nitric acid.
On heating it tends to decompose into a red gas, nitrogen
peroxide, and free oxygen.

2N2O6 = 4NO2 + O2

Nitrogen peroxide, NO2 or N2O4

Nitrogen peroxide is frequently produced when we heat the
nitrates of the heavy metals. The substance which we

158 INORGANIC CHEMISTRY

generally employ in the laboratory is lead nitrate, which on
heating decomposes according to the equation —

2Pb(N08)2 = 2PbO + 4NO2 + O2

a mixture of nitrogen peroxide and oxygen coming off. The
nitrogen peroxide may be condensed to a light brown liquid,
by passing it through a tube immersed in a freezing mixture
of ice and salt. The oxygen is not condensed to a liquid
under these circumstances, and so the two substances may
be separated.

Nitrogen peroxide can also be easily produced by the direct
union of nitric oxide and oxygen, in accordance with the
equation —

2NO + 02 = 2NO2

Nitric oxide Oxygen Nitrogen peroxide

This union takes place at the ordinary temperature, and is
rendered evident by the fact that nitric oxide and oxygen
are both colourless gases, whilst nitrogen peroxide is a reddish
brown gas.

If nitrogen peroxide is heated, the colour may be seen to
darken until the gas becomes practically opaque. On cooling
again, the colour falls off to its original tint. These colour
changes are connected with the following action : —

N2O4 <t 2NO

This action is a reversible action, decomposition taking place
on heating, and recombination taking place on cooling. A
decomposition of this kind is called a dissociation, the essential
character of dissociation being reversibility and the production
of a larger volume of gas by the decomposition.

At ordinary temperatures, nitrogen peroxide consists partly
of the gas NO 2, and partly of the gas N2O4, the proportions
of these two substances which are present being variable with
the temperature, the NO 2 increasing in amount as the
temperature is raised. The N2O4 has a somewhat pale
reddish brown colour, while the NO 2 has a very deep and
almost opaque brown colour.

Nitrogen peroxide at once dissolves in water, but it does
not dissolve as such, the water decomposing it with formation

NITROGEN 159

of other compounds of nitrogen. If the water is cold and
in small quantity, the action takes place as follows : —

2NO2 + H2O = HNO3 + HNO2

Nitric acid Nitrous acid

With a large quantity of water at a higher temperature,
nitric oxide is produced, owing to the decomposition of the
nitrous acid ; and the whole action may be represented thus :

3NO2 + H2O = 2HNO8 + NO

Nitric acid Nitric oxide

Nitrogen triozide,. N2O3

This oxide of nitrogen is very unstable. It can easily be
prepared by mixing nitrogen peroxide with nitric oxide, the
partial combination which takes place being represented by
the equation —

NO2 + NO = N2O3

Nitrogen peroxide Nitric oxide Nitrogen trioxide

If the oxides be taken in the proper proportions, and the
gases liquefied by means of a freezing mixture, the liquid
obtained differs altogether in colour from the brown nitrogen
peroxide, being of a deep greenish blue. Probably this liquid
contains the substance N2O3, but when i it is converted into
gas, the nitrogen trioxide decomposes to a very great extent
into nitric oxide and nitrogen peroxide. We have thus the
balanced action —

N2O3 <^ NO2 + NO

which is comparable to the dissociation of nitrogen peroxide
by heat —

N2O4 ^ NO2 + NO

Nitrogen trioxide is most conveniently prepared by heating
nitric acid with arsenious oxide (white arsenic), which reduces
the nitric acid with formation of nitrogen trioxide.

i6o INORGANIC CHEMISTRY

When nitrogen trioxide is dissolved in cold water, it yields
a bluish liquid which contains nitrous acid.

N2O3 + H2O = 2HNO2

Nitrogen trioxide Water Nitrous acid

Nitrogen trioxide would thus seem to be the anhydride of
nitrous acid. Nitrous acid, however, has not been obtained
in the pure state, and the blue solution decomposes readily
on heating, with formation of nitric acid, which remains dis-
solved, and nitric oxide, which escapes as a gas.

3HNO2 = HNO3 + .2NO + H2O

Nitrous acid Nitric acid Nitric oxide

Nitric oxide, NO

Nitric oxide is one of the commonest products of the action
of nitric acid on a metal. Thus, when lead, copper, or silver
is attacked by nitric acid, nitric oxide is produced, according
to the equations —

3Pb + 8HNO3 = 3Pb(N03)2 + 2NO + 4H2O
3Cu + 8HNO3 = 3Cu(N03)2 + 2NO + 4H2O
3Ag + 4HNO3 = 3AgN03 + NO + 2H2O

In the laboratory, it is generally produced by the action of
slightly diluted nitric acid on copper turnings. A convenient
apparatus for the production of nitric oxide on the small
scale from these materials is shown in fig. 30.

Nitric oxide is a colourless gas which is incompressible to
a liquid at the ordinary temperature, and is only very slightly
soluble in water. As soon as it comes into contact with
oxygen or air, it combines with the oxygen, according to
the equation —

2NO + 02 = 2NO2

and forms red fumes of nitrogen peroxide. It will not support
the combustion of a taper, but if burning phosphorus is
brought into contact with it, the phosphorus continues to
burn, decomposing the nitric oxide by combining with the
oxygen and liberating the nitrogen.

4P + loNO = 2P2O5 + 5N2

Nitric oxide gives a characteristic dark brown coloration

NITROGEN i6r

with a solution of ferrous sulphate, the production of which
is a convenient means for detecting the presence of nitric
acid, or of a nitrate, in solution.

The solution suspected to contain
a nitrate is mixed with a little fer-
rous sulphate solution, and this mixed
solution is poured carefully down the
side of a test tube, whitih contains
a little concentrated sulphuric acid,
care being taken that the two' liquids
are, not allowed to mix. If a nitrate
is present, a brown ring is produced
at the surface of contact of the two
liquids. The nitric acid formed at
the surface from the nitrate and the
strong sulphuric acid is reduced by
the ferrous sulphate to nitric oxide,
which dissolves in the excess of the
ferrous sulphate with production of
the brown compound. As this brown
compound is easily destroyed by
heat, the sulphuric acid must not *'*^- ^NMro^l"°° *"
be allowed to mix in bulk with the copper imningt are con-
aqueous solution, for then so much lained in Ihe uppa- p«l of Ihe

heal would be evolved as to destroy si^irify"'d[iut^'"'wLib"^tKr,'
the brown ring and thus render the ^ "•"] ™ '''^j"'n'i„|J'oJi'|^

test useless. CKap;> Ihtough tbe delivEiy

lalKZ>.

Nitrous oxide, N3O
We have seen that when ammonium nitrite is heated it decom-
poses into water and nitrogen, according to the equation —

NH.NOV = Ns -I- 2HsO
Now ammonium nitiate contains more oxygen than ammonium
nitri'/e, and when heated decomposes in a similar manner, with
production, however, not of nitrogen, but of nitrogen mon-
oxide or nitrous oxide.

NH.NO3 = N2O + aHjO

Ammonium nitrate Nitrous oxide

Nitrous oxide, like nitric oxide, is a colourless gas. It

l62

INORGANIC CHEMISTRY

possesses a faint sweetish smell and taste, and is moderately
soluble in cold water.

It may also be prepared by heating a mixture of ammonium
sulphate and potassium nitrate, the action which takes place
being represented by the equation —

(NH4)2S04 + 2KNO8 = 2N2O + 4H2O + K2SO4

A little nitric oxide and nitrogen peroxide are formed at the
same time. These, however, can easily be removed by
passing the gas through ferrous sulphate solution. Since
the gas is soluble in cold water, it is usually collected over
warm water, in which it is much less soluble.

Fig. 31. — Preparation of Nitrous Oxide.

The gas produced in the flask A is washed by bubbling through ferrous sulphate
solution contained in the wash>bottIe fF, and collected over warm water.

Nitrous oxide -is sometimes known by the name of laughing
gas. When inhaled it produces a sense of exhilaration and
hysterical laughter, which is afterwards succeeded by com-
plete insensibility. Nitrous oxide is therefore used as an
anaesthetic for the smaller operations of dentistry and the
like. Corresponding to its solubility in water, we find that it
can be compressed to a liquid at the ordinary temperature,
and it is usually supplied to dentists in this form, the liquid
being contained in small steel bottles.

Nitrous oxide can easily be distinguished from nitric oxide
owing to the fact that it does not combine with oxygen to
produce the red nitrogen peroxide, and that it is a vigorous
supporter of combustion. When a lighted taper is introduced
into a jar of nitrous oxide it burns with increased brilliancy,

NITROGEN 163

and when a glowing splinter of wood is brought into contact
with the gas it bursts into flame just as if it had been placed
in a jar of oxygen.

Nitrous oxide is thus a better supporter of combustion than
nitric oxide, although it contains proportionately less oxygen,
and either of the gases is a better supporter of combustion
than nitrogen peroxide, which contains more oxygen still.
We therefore see that the power of supporting combustion
depends, in the case of a compound, not so much on the
quantity of oxygen it contains, but on the ease with which
it parts with its oxygen. Nitrous oxide, which has a com-
paratively small amount of oxygen, parts with the oxygen it
has very readily, and is thus a good supporter of combustion.

It is necessary, however, that the substance which is to
burn in the nitrous oxide should be at such a temperature as
to decompose the nitrous oxide and abstract the oxygen from
the nitrogen. If we take some sulphur which is only burning
feebly, and place it in a jar of nitrous oxide, the sulphur will
be extinguished. If, on the other hand, the sulphur is
burning brightly before it is placed in the nitrous oxide, it
will now burn more brightly than ever. We have thus the
facts that feebly burning sulphur is extinguished by nitrous
oxide, and brightly burning sulphur burns in the nitrous
oxide with increased brilliancy. In the first case, the tem-
perature of the burning sulphur is not sufficient to enable
it to abstract oxygen from the nitrous oxide, and in the second
case the temperature is sufficiently high to effect this decom-
position, and the proportion of oxygen which is thus produced
is greater than the proportion of oxygen in air, and thus the
brilliancy of the combustion is increased.

Nitrous acid a>nd the Nitrites

When nitrates of most metals are heated, we find that the
residue consists of an oxide of the metal. If potassium or
sodium nitrate, however, is heated, the residue consists of
potassium or sodium nitrite^ the decomposition taking place
according to the equations —

2KNO8 = 2KNO2 + O2

2NaN03 = 2NaN02 + O2

These nitrites are solid substances, like the nitrates. They

1 64 INORGANIC CHEMISTRY

give with ferrous sulphate and sulphuric acid the ring test in
much the same way as the nitrates. They can easily be dis-
tinguished from the nitrates, however, by the simple addition
of a dilute acid such as sulphuric acid or hydrochloric acid.
If they are gently warmed with these acids, copious brown
fumes are evolved, the nitrites being decomposed with for-
mation of nitrous acid, which at once decomposes with
formation of nitric oxide and nitrogen peroxide. Nitrates
give no red fumes under similar conditions.

The nitrites, when acidified, can act either as oxidising
agents or as reducing agents. When they act as reducing
agents, the nitrous acid takes up oxygen and is converted into
nitric acid. When they act as oxidising agents, the nitrous
acid gives up oxygen and is converted into water and nitric
oxide, the equations being as follows : —

HNO2 + (O) = HNOa

HNO2 + (H) = NO + H2O

Here, again, we see that oxidising power depends not so much
on the quantity of oxygen, as on the ease with which the
oxygen is given up. A cold dilute solution of nitric acid
HNO3 is by no means a powerful oxidising agent, although
it contains plenty of oxygen, while a cold dilute solution of
nitrous acid HNO2 will part with oxygen at once to become
nitric oxide and water.

Nitric acid

The preparation of nitric acid, or aqua fortiSy by the action
of sulphuric acid on a nitrate (usually sodium nitrate) has been
described on p. 47. Nitric acid differs from the other common
mineral acids principally in being a very powerful oxidising
agent when undiluted with water. Thus when brought into
contact with powdered charcoal the concentrated acid inflames
it at once with production of carbon dioxide. The nitric acid
in such reactions is itself usually reduced to an oxide of
nitrogen, but sometimes to nitrogen or to ammonia. Examples
of its oxidising action, by itself and in solution, are given on
pp. 104-106, 157-164, 187, 212.

CHAPTER XXIV

HTDBOGEN

Water is practically the only source of hydrogen and its
compounds. Water can be split up into its elements — namely,
hydrogen and oxygen — directly by means of the electric
current. In order, however, to make water conduct electri-
city, sulphuric acid in small quantity must be added to it
before the current can be passed. Oxygen appears at the
positive pole, and hydrogen appears at the negative pole, the
two being in the proportions necessary to form water. By
weight these proportions are roughly, i part of hydrogen to
8 parts of oxygen ; by volume they are 2 measures of hydrogen
to I measure of oxygen (fig. 20).

Hydrogen can also be liberated from water by acting upon
the water with various substances which are capable of
removing oxygen. Thus, if the metal sodium is thrown into
water a vigorous action at once takes place, with production of
sodium hydroxide and hydrogen, according to the equation —

2Na + 2H^0 = 2NaOH + H

Other metals, such as zinc, iron, and aluminium have no
action upon water at the ordinary temperature; but if they
are raised to a red heat, and if steam is then passed over
them, they decompose the steam with formation of an oxide
of the metal and hydrogen gas, the equations being —

H2O

FeO

H2O

ZnO

2AI

3H2O

Al.Oa

3H.

Some metals, such as mercury, silver, and gold, are in-
capable of decomposing water at any temperature (p. 79).

It frequently happens that two metals which cannot decom-
pose water singly can decompose water when in contact

165

1 66 INORGANIC CHEMISTRY

with it together. Thus neither mercury nor aluminium can
decompose water at the ordinary temperature. Yet, if
aluminium is coated with mercury, the aluminium-mercury
couple^ as it is called, can decompose water easily at the
ordinary temperature, and very rapidly at the boiling point,
aluminium hydroxide being produced. The explanation of this
is that a difference of electric state is set up between the two
metals, which is sufficient to decompose the water. Similarly,
zinc when coated with copper can decompose water with
production of zinc hydroxide, although metallic zinc alone
has no action on water below a red heat.

Carbon also is capable of decomposing water when heated
to a high temperature, the product of the reaction being
hydrogen and carbon monoxide (p. 149).

C + H2O = CO + H2

The ordinary method of preparing hydrogen in the labora-
tory is by the action of hydrochloric or dilute sulphuric acid
on a metal, the metal usually chosen being zinc, and a Kipp
machine being commonly used as the generator.

Zn + H2SO4 = ZnS04 + Hg

Hydrogen prepared by any of these methods is a colourless
and odourless gas. Frequently, however, the crude gas has a
distinctly unpleasant smell, which is not the smell of the
hydrogen, but of some impurity derived from the materials
used in its production. Hydrogen is singular in many re-
spects. It is the lightest gas known, and one of the most
difficult to condense to a liquid, an extremely low degree
of temperature being necessary to effect liquefaction. The
temperature of boiling hydrogen is - 253** — that is, only 20'
above the absolute zero of temperature. At this temperature
practically all other gases are condensed to solids. In accord-
ance with this very low boiling point, we find that the gas
is very slightly soluble in water, the solubility being approxim-
ately the same as that of nitrogen — namely, 0.02.

Hydrogen has also the smallest combining weight of any
of the elements, and all the other combining weights are for
ordinary purposes referred to it as unity. If we weigh the
amount of hydrogen which occupies 22.4 litres at normal

HYDROGEN 167

temperature and pressure we find that it is two grams. The
molecular weight is therefore 2 and the molecular formula Hg.

Since hydrogen is the lightest of all gases, the densities of
other gases are very frequently referred to the density of
hydrogen under the same conditions as unity. It therefore
follows that the molecular weights of gases or vapours are equal
to double their densities when these densities are referred to
hydrogen, for the density of hydrogen is chosen equal to i
whilst its molecular weight is equal to 2. To get the molecular
weight of any gas, then, we have simply to multiply its
density referred to hydrogen by 2.

Thus oxygen is found to be sixteen times heavier than an
equal bulk of hydrogen when the two gases are measured at
the same temperature and pressure. According to the above
rule, oxygen will therefore have the molecular weight 2 x 16
= 32. We can easily see that this is in accordance with Avo-
gadro's principle (p. 1 14), for if 2 grams of hydrogen go into
22.4 litres at normal temperature and pressure, sixteen times
as much as this of oxygen will go into the same space under
the same conditions. That is, 32 grams of oxygen will be
contained in 22.4 litres at N.T.P., or the molecular weight of
oxygen is 32.

Hydrogen is a combustible gas which bums in air or oxygen
with a non-luminous flame, the only product of the com-
bustion being water, which is formed according to the
equation —

2H2 + 02 = 2H2O

2 vols. I vol.

A mixture of hydrogen and oxygen in the above proportions
by volume explodes with great violence when a light is applied
to it, and is often called detonating mixture. It can be made
with the components in the proper proportions by the elec-
trolysis of acidulated water.

Hydrogen, as we have seen, acts as a reducing agent, inas-
much as it is capable of removing oxygen from many oxygen
compounds when these are heated in a stream of hydrogen
gas.

Coal-gas contains 40 to 50 per cent, by volume of hydro-
gen, which gives out a great deal of heat when burned, but
no light. On account of this amount of hydrogen which

i68 INORGANIC CHEMISTRY

it contains, coal-gas can sometimes be used in the laboratory
as a gaseous reducing agent. Thus, if it is passed over lead
oxide, the oxygen will be removed by the hydrogen, and
metallic lead will be left.

Hydrogen peroxide, H2O2

Water is not the only substance formed by the union of
oxygen and hydrogen, although when these two elements
combine directly, water is the sole product. There is another
compound of hydrogen and oxygen — hydrogen peroxide,
which for a given amount of hydrogen contains twice as much
oxygen as water. Hydrogen peroxide has utterly different pro-
perties from water, being in its nature a very weak acid, and
also a powerful oxidising agent.

Salts of hydrogen peroxide can easily be prepared. Thus
if we heat metallic sodium in air, it burns to produce sodium
peroxide Na202, which is the sodium salt of hydrogen per-
oxide. When therefore we treat this sodium salt of the weak
acid, hydrogen peroxide, with a stronger acid, the stronger
acid turns out the hydrogen peroxide and takes the base.
Thus, sodium peroxide and sulphuric acid give in aqueous
solution, sodium sulphate and hydrogen peroxide —

Na202 + H2SO4 = Na2S04 + H2O

Barium peroxide is a similar salt of hydrogen peroxide. It
can be formed by heating barium monoxide carefully in air
(p. 128). When mixed with water, in which it is only slightly
soluble, and treated with an equivalent quantity of sulphuric
4cid, insoluble barium sulphate is produced, and hydrogen
peroxide passes into solution. The barium peroxide may
also be decomposed by means of a current of carbon dioxide
according to the following equation : —

Ba02 + CO2 + H2O - BaCOg + H2O2

Barium peroxide Hydrogen peroxide

The barium carbonate is insoluble, and may be easily separ-
ated from the solution of hydrogen peroxide.

The solution of hydrogen peroxide prepared in any of these
ways decomposes, on heating, into water and oxygen, which

HYDROGEN 169

comes off in the form of bubbles of gas, the equation for the
decomposition being —

2H2O2 = 2H2O + O2

At the ordinary temperature, hydrogen peroxide in solution
is moderately stable, especially when the solution is dilute.
If the solution is kept over concentrated sulphuric acid in
a vessel which has been exhausted of air, the water of the
solution evaporates, and is absorbed by the sulphuric acid,
leaving behind a syrupy residue of pure hydrogen peroxide,
which is almost half as heavy again as water. Pure hydrogen
peroxide is stable if the temperature remains in the neighbour-
hood of the freezing point, but on a warm day it decomposes
with evolution of bubbles of oxygen, and at the boiling point
it evolves oxygen with explosive violence. Some substances,
such, as silver, cause it to evolve oxygen very vigorously at the
ordinary temperature by mere contact, the substances them-
selves apparently remaining unchanged.

When hydrogen peroxide is brought into contact with silver
oxide, the following action occurs :: —

AggO + H2O2 = 2Ag -I- O2 + H2O

Silver oxide Silver

Here the hydrogen peroxide apparently behaves, as a reduc-
ing agent so far as the silver oxide is concerned. With
powerful oxidising agents, it often presents this character.

When a few. drops of potassium bichromate solution, acidi-
fied with acetic acid, is added to a solution of hydrogen per-
oxide, an intense blue colour is produced. This blue colour
is attributed to perchromic acid, and its production can be
made use of as a test both for hydrogen peroxide, and for
chromates. The hydrogen peroxide in this reaction behaves
as an oxidising agent, giving up oxygen to the bichromate,
and being itself reduced to water.

Hydrogen peroxide, therefore, like nitrous acid, possesses
both oxidising and reducing properties, either of which may
be developed according to circumstances.

CHAPTER XXV

CHLOBINE

Common salt, or sodium chloride^ NaCl, is the universal source
of chlorine and all its compounds. This substance occurs not
only in enormous quantity in sea water, from which it may be
obtained by evaporation, but also in brine springs and as solid
rock salt.

Chlorine is now obtained from sodium chloride on the
manufacturing scale by electrolytic decomposition. The negative
radical of the chloride travels to the positive electrode, which
consists of carbon, and is there discharged with formation of
free chlorine.

2CI = CI2 .

Chloride radical Chlorine gas

Another manufacturing process for the production of chlorine,
which can also be used in the laboratory, is the dehydrogenisa-
tion of hydrochloric acid. This acid when treated with many
oxidising agents loses hydrogen in the form of water, the
chlorine being at the same time liberated as gas. The oxidis-
ing agent most commonly employed is manganese dioxide, and
the action may be represented by the equation —

MnOg + 4HCI = MnCla + 2H2O + Clg
Manganese dioxide Manganese chloride Chlorine

In the laboratory the black manganese dioxide is warmed in
a flask with a moderately concentrated solution of hydro-
chloric acid, and the chlorine gas, which is evolved in a
steady stream, may be dried by passing through strong sul-
phuric acid. The gas is generally collected by downward
displacement, since it is somewhat soluble in water and
readily attacks mercury.

Instead of using hydrochloric acid, we may substitute a

170

CHLORINE 171

mixture of sodium chloride and strong sulphuric acid which
is capable of producing it. The action is then represented by
the equation —

MnOa + 2NaCl + 2H2SO4 = MnS04 + Na2S04 + 2H2O + Clg

Manganese Manganese Chlorine

dioxide sulphate

Chlorine is a greenish yellow gas, with a pungent suffocating
odour, by means of which.it may be easily recognised. It
is moderately soluble in water, the solution being termed
chlorine water^ and can be condensed to a liquid at the
ordinary temperature by means of pressure. Its molecular
weight is 71, so that the gas is nearly two and a half times as
heavy as air under the same conditions.

Chlorine is a very active element, readily attacking metals
with formation of metallic chlorides, and uniting directly with
such non-metals as hydrogen, sulphur, and phosphorus. Thus
phosphorus, if brought into a jar of chlorine, takes fire spon-
taneously with production of the chlorides of phosphorus.
Chlorine has a special attraction for hydrogen, not only uniting
with it directly, but even removing it from compounds in
which it is a constituent, hydrochloric acid being in each case
produced by the union. Thus chlorine comes to be used as
a dehydrogenising (or oxidising) agent, its practical application
being mostly for bleaching or for disinfecting. On account
of its gaseous nature and insupportable odour, chlorine is
very seldom used in the free state, but is at once converted
into bleaching powdery from which chlorine can at any time be
readily obtained, and which possesses both the bleaching and
disinfecting properties of the element itself.

The dehydrogenising action of chlorine may readily be seen
by placing a piece of filter paper saturated with turpentine in
a vessel filled with chlorine. Turpentine is a compound of
carbon and hydrogen, and from it chlorine abstracts the
hydrogen with such vigour that the liquid bursts into
flame, with formation of a dense black smoke consisting of
particles of carbon, on which chlorine has no action. For the
same reason, if a taper or a jet of coal-gas is burnt in a jar
of chlorine, the flame produced is exceedingly smoky owing to
the separation of carbon.

The bleaching properties of chlorine may be seen by adding

172 INORGANIC CHEMISTRY

a little chlorine water to ink, or a solution of indigo, the
colour of these substances being at once discharged.

Hydrochloric acid, HOI

Hydrochloric acid is prepared both in the laboratory and on
the large scale by heating common salt with strong sulphuric
acid. The first action, which occurs without special heating,
is the production of hydrochloric acid and acid sodium
sulphate.

NaCl + H2SO4 = HCl + NaHS04

The hydrochloric acid is a gas, and so escapes; the acid
sodium sulphate is a solid, and remains behind (fig. 17). If
more salt is now added and the temperature raised, a second
action takes place — viz.

NaCl + NaHS04 = HCl + Na2S04

more hydrochloric acid and normal sodium sulphate being pro-
duced. These two actions may both be expressed in one
equation, derived from the preceding equations by adding
their corresponding sides together, and eliminating terms
common to both sides. The equation for the whole action
is then —

2NaCl + H2SO4 = 2HCI + Na2S04

As a rule, in the laboratory the action is not pushed beyond
the first stage.

Hydrochloric acid is very seldom used in the form of gas.
It is almost invariably dissolved in water, in which it is exces-
sively soluble. What is known as strong or concentrated
hydrochloric acid is an aqueous solution containing about
one-third of its weight of the pure acid. The commercial
solution is sometimes known under the old names of spirit of
salt^ or muriatic acidy derived from the Latin, muria^ brine.

Hydrochloric acid is a colourless gas with a very pungent
odour, quite distinct, however, from that of chlorine. Its
molecular weight is 36.5, corresponding to the formula HCl,
so that the gas is considerably heavier than air, and may be
collected by downward displacement. It is very soluble in
water, one volume of which under ordinary conditions dis-

CHLORINE 173

solves nearly 500 volumes of the gas. From this solution
hydrochloric acid (unlike ammonia, p. 156) cannot be expelled
by boiling, as after a time the hydrochloric acid and water boil
off together in constant proportions until all the liquid has
disappeared. When brought into contact with air, it fumes
strongly, the white fumes consisting of small particles of liquid
solution of hydrochloric acid produced by the condensation of
the gas with the moisture of the air. The gas may be liquefied
at the ordinary temperature by the application of pressure.

As has already been indicated, hydrochloric acid may be
produced by the direct union of hydrogen and chlorine.
These gases can be mixed and kept for an indefinite time
in the dark without undergoing alteration, but if the mixture
is exposed to light, a slow union goes on with production of
hydrochloric acid, according to the equation —

H2 + CI2 = 2HCI

I vol. I vol. 2 vols.

Should the hydrogen and chlorine be present in the mixture
in exactly the proportions required for the reaction — ue, in equal
volumes — exposure to a bright light will cause the union to
take place so rapidly as to be explosive. Direct sunlight or
the light from burning magnesium ribbon is usually sufficient
to determine the explosive combination of the two gases. In
all cases the sudden combination may be brought about by
applying a lighted taper directly to the mixture.

Many metals when heated in an atmosphere of hydrochloric
acid are converted into chlorides, hydrogen being liberated at
the same time. Ferrous chloride, for example, may be pre-
pared in this way, the action taking place according to the
equation —

Fe + 2HCI = FeClg + Hg

Iron 2 vols. Ferrous chloride I vol.

The volume of hydrogen liberated is always equal to half the
volume of the hydrochloric acid gas which has been decom-
posed. It may be said in general, that the metals which are
acted on by dilute hydrochloric acid solution, also decompose
the gas, either at the ordinary temperature or when heated.
Hydrochloric acid, being a strong acid, readily attacks most

174 INORGANIC CHEMISTRY

basic oxides or hydroxides, with formation of a metallic chloride
and water, thus —

FeO + 2HCI = FeClg + H2O
F^(0H)2 + 2HCI = FeCla + 2H2O

In one point hydrochloric acid differs from the other strong
common mineral acids, nitric and sulphuric acids, — it never
acts as an oxidising agent. We therefore find that its action
on metals is more restricted than the action of these other
acids : it closely resembles the action of dilute sulphuric acid,
which likewise possesses no oxidising properties.

Oxides of Chlorine

GMorine monoxide, CI2O. — Chlorine and oxygen will not,
under any circumstances, combine directly. They may be
made to unite, however, if the oxygen is previously combined
with some other element. In this way chlorine monoxide
may be formed by passing a current of chlorine through a long
tube containing dry mercuric oxide. The chlorine unites with
the mercury and with the oxygen simultaneously, forming
mercuric chloride and chlorine monoxide, thus —

HgO + 2CI2 = HgCl2 + CI2O

If the mercuric oxide is in excess, a compound HgO,HgCl2
intermediate between the oxide and chloride is produced * —

2HgO + 2CI2 = HgO,HgCl2 + CI2O

Basic mercuric
chloride

Chlorine monoxide is a dense yellow gas at the ordinary
temperature, which even on gentle heating decomposes into
its elements with explosive violence. Care must therefore be
taken in its preparation.

* According to the definition given on p. 65, basic mercuric chloride
should have the formula Hg(OH)Cl or Hg(OH)a,HgCla, and be inter-
mediate in composition between the base HgfOH)^ and the normal salt
HgClj. The name basic salt is frequently given, however, to compounds
intermediate between the basic oxide and the normal salt. Thus the com-
pound HgjOClj, or HgO,HgClj is termed basic mercuric chloride.

CHLORINE 175

It dissolves readily in water, forming a solution of hypo-
chlorous acid HCIO, and is therefore frequently termed
hypochlorous anhydride.

CI2O + H2O = 2HCIO

Chlorine monoxide Hypochlorous acid

Chlorine peroxide, CIO 2. — This substance resembles chlorine
monoxide in many ways, and is more frequently met with,
since its production from a chlorate and strong sulphuric acid
is generally employed as a test for chlorates. Its forma-
tion from potassium chlorate is represented by the following
equation : —

3KCIO8 + H2SO4 = KCI94 + K2SO4 + 2CIO2 + H2O

Potassium Potassium Chlorine

chlorate perchlorate peroxide

The gas has a bright yellow colour, not greenish yellow like
chlorine, and a characteristic odour, which renders it easy of
detection. At a temperature considerably below the boiling
point of water, it decomposes explosively into its elements. At
the ordinary temperature, combustible substances like sulphur
catch fire when brought into contact with it.

Action of Chlorine on Alkalies

When chlorine is led into a cold solution of caustic soda,
caustic potash, or calcium hydroxide, it forms with these sub-
stances bleaching solutions which contain the chloride and the
hypochlorite of the metal.

2NaOH

+ CI2 = NaCl + NaOCl +

H2O

2KOH

+ CI2 = KCl + KOCl +

H2O

2Ca(OH)2

+ 2CI2 = CaClg + Ca(0Cl)2 +

Chloride Hypochlorite.

2H2O

The bleaching action of the solutions prepared in this way is
due to the hypochlorites^ the chlorides of the alkalies having no
bleaching properties.

If the alkaline solution is hot instead of cold, the solution
obtained by the action of chlorine does not bleach, and con-

176 INORGANIC CHEMISTRY

tains clilorate instead of hypochlorite. The equations repre-
senting this type of action are as follows : —

6NaOH + 3CI2 = sNaCl + NaClOa + 3H2O
6K0H + 3CI2 = 5KCI + KCIO3 + 3H2O
6Ca(0H)a + 6CI2 = sCaClg + Ca(C103)2 + 6H2O

Chloride Chlorate.

The production of chlorate instead of hypochlorite at the
higher temperature depends on the fact that when a solution of
a hypochlorite is boiled it becomes converted into a mixture of
chloride and chlorate, thus —

3NaOCl = 2NaCl + NaClOs

Hypochlorite Chloride Chlorate

If chlorine is permitted to act on dry calcium hydroxide
instead of on a mixture of calcium hydroxide and water, a very
important substance is produced — namely, blea.cliiiig powder.
This substance is, as it were, intermediate between calcium
chloride and calcium hypochlorite, and when dissolved in
water gives a bleaching solution which contains both of these
salts. The actions may be represented by the equations —

Ca(0H)2 + CI2 = CaOCl2 + H2O .

Bleaching powder.

2CaOCl2 = CaCl2 + Ca(0Cl)2

Bleaching powder Chloride Hypochlorite

Bleaching powder, however, does not consist entirely of a
compound CaOCl2, as a portion of the solid calcium hydroxide
is always unattacked by the chlorine. Thus, when bleaching
powder is treated with water so as to form a bleaching solution,
a quantity of the sparingly soluble calcium hydroxide remains
undissolved.

The action of chlorine on ammonia differs entirely from its
action on other alkalies* In this case no hypochlorite or
chlorate is formed, the ammonia being decomposed with
liberation of nitrogen. This action occurs whether the
ammonia is present in the gaseous state or in aqueous
solution, and may be represented by the equation —

2NH8 + 3CI2 = N2 + 6HC1
If the ammonia is present in excess, the hydrochloric acid is,

CHLORINE 177

of course, converted into ammonium chloride. If the chloride
is present iii excess, the highly explosive nitrogen trichloride
may be formed, thus —

NH3 + 3CI2 = NCI3 + 3HCI

Nitrogen trichloride.

The action of chlorine on the basic mercuric oxide has
already been alluded to on p. 174.

Bleaching Action of Chlorine and Hypochlorites

That the bleaching action of chlorine is sometimes a process
of oxidation and not of direct dehydrogenisation (compare
p. 99) may be illustrated by the following experiment : — If a
piece of carefully dried red calico is placed in a jar of dry
chlorine, there is scarcely any bleaching action ; but if the
experiment is repeated with a piece of moist calico, the bleach-
ing action is almost instantaneous. The chlorine, therefore, in
absence of water, is unable to destroy the colouring matter by
its own dehydrogenising action. If water is present it is
simultaneously attacked by the chlorine and by the colouring
matter, the chlorine taking the hydrogen of the water, and the
colouring matter the oxygen, thus —

H2O + CI2 = 2HCI + O

The oxygen represented in the equation goes to oxidise the
colouring matter to a colourless substance.

The action of hypochlorous acid is closely similar. It gives
up oxygen to the colouring matter, and is converted into hydro-
chloric acid, thus — '

HOCl = HCl + O

It should be noticed that a given amount of chlorine in* hypo-
chlorous acid yields as much bleaching oxygen as twice that
amount of free chlorine.

When hydrochloric acid is added to a solution of hypo-
chlorous acid, chlorine and water are produced, in accordance
with the equation —

HOCl + HCl = CI2 + H2O
The bleaching solutions, which contain both chloride and

178 INORGANIC CHEMISTRY

hypochlorite, are not themselves powerful bleaching agents
Hke hypochlorous acid, or chlorine and water, since they do
not so readily part with oxygen. If these solutions are
acidified, however, hypochlorous acid is liberated, and this at
once acts as a bleaching agent. Carbon dioxide, although the
anhydride of a very weak acid, is still strong enough to liberate
hypochlorous acid with ease from a hypochlorite, so that if the
substance to be bleached is soaked in a bleaching solution and
exposed to the air, the carbonic acid present in small quanti-
ties in the atmosphere will slowly liberate hypochlorous acid
and effect the bleaching action.

When an acid is added to a bleaching solution, hydrochloric
acid is set free along with hypochlorous acid, and that the
more readily, as the foreign acid is stronger. But we have
already seen that these two acids cannot exist together in
solution, reacting together to form water and chlorine. We
therefore get chlorine as a product when a bleaching solution
is strongly acidified. The smell of bleaching powder and of
bleaching solutions in general is due to hypochlorous acid,
liberated in small quantity from the hypochlorite by the action
of atmospheric carbonic acid.

The bleaching solution from bleaching powder in no way
differs, so far as its bleaching properties are concerned, from
the other bleaching solutions, since it behaves precisely as a
mixture of calcium chloride and calcium hypochlorite. The
yarn or other substance to be bleached is generally freed
from grease, etc., by boiling in a weak alkaline bath of wash-
ing soda, and then dipped alternately into weak solutions of
bleaching powder and of sulphuric acid. The sulphuric acid
acts on the hypochlorite, liberating hypochlorous acid and
chlorine, which then perform the work of bleaching.

The reason why bleaching powder is used so extensively in
place of the bleaching solutions lies in the fact of its being
a solid, which can be kept for a long time without much
alteration. It can, therefore, be conveniently stored and
transported, which is not the case with bleaching solutions.

Chlorates and Perchlorates

Chlorates do not occur in nature, but, as we have seen, are
prepared by the action of chlorine on alkaline solutions in the

CHLORINE 179

heat. The chlorate which is most frequently used is potas-
sium cMorate, KCIO3. This substance might be made
directly from caustic potash and chlorine, according to the
equation —

6K0H + 3CI2 = 5KCI + KCIO3 + 3H2O

but on the large scale this mode of preparation is needlessly
expensive, on account of the quantity of comparatively dear
potassium hydroxide which is converted into the compara-
tively cheap potassium chloride. The inexpensive alkali,
calcium hydroxide, is therefore substituted for potassium
hydroxide, so that calcium chlorate and calcium chloride are
produced by the action of the chlorine; and then a strong
solution of the calcium chlorate is mixed with a strong solution
of potassium chloride. On cooling, the sparingly soluble
potassium chlorate separates out of the solution, the chemical
action being —

Ca(C103)2 + 2KCI = CaClg + 2KCIO3

It is chiefly on account of the ease with which potassium
chlorate can be purified by crystallisation that it is prepared on
a large scale in preference to other chlorates.

Chlorates, like nitrates, part readily with their oxygen when
heated, and are consequently used as sources of oxygen in the
production of fireworks, compositions for the tips of matches,
etc. Such mixtures consist essentially of a combustible sub-
stance {e.g, sulphur, sugar, charcoal), and an oxidising substance
{e,g, nitrate or chlorate). When the mixture is heated to a
certain temperature, chemical action begins, the combustible
substances being oxidised by the nitrate or chlorate.

In neutral aqueous solution chlorates, like nitrates, do not
behave as strong oxidising agents, but, on the addition of acid,
the oxidising properties become apparent. This is due to the
formation of the strongly oxidising chloric acid itself, and also
of chlorine peroxide and free chlorine. Thus, when sulphuric
acid is added to a chlorate, chlorine peroxide is produced,
according to the equation —

3KCIO3 + H2SO4 = KCIO4 + K2SO4 + 2CIO2 + H2O
Hydrochloric acid generates by its action on a chlorate a

::J^>*==^_

-t»y

,4^'

SI%.

rL5»-*-'

.- *'

r ' ;• ^5> ^ ^S>^

I*-

s- X

^ '-C***^"
?iW -^^

^^. -

-Vv

n*^ c =^r

J-?v

>%v^

CHLORINE i8i

general action. On further heating, the perchlorate decom-
poses into chloride and oxygen, thus —

KCIO4 = KCl + 2O

2 i

so that if the chlorate is heated to a sufficiently high tempera-
ture, chloride and oxygen are the only products.

Most chlorates, when heated, decompose like potassium
chlorate, but some give off chlorine as well as oxygen.

All the chlorates and perchlorates are soluble in water, the
potassium salts being amongst the most sparingly soluble.

Chloric acid, HCIO3, and perchloric acid, HCIO4, are both
liquid substances, and very strong oxidisers. If dropped in the
pure state on wood, paper, charcoal, or other organic material,
oxidation at once sets in, sometimes with explosive violence.

The perchlorates are easily distinguished from the chlorates
-by their not giving chlorine peroxide on treatment with sul-
phuric acid.

Nitrogen trichloride

Chlorine will not combine directly with nitrogen, but will
do so when it has the opportunity of combining with hydrogen
simultaneously. Thus it has already been stated that nitrogen
trichloride is formed by the action of chlorine on ammonia.
The best method of preparing nitrogen trichloride, however, is
by the action of chlorine on a saturated solution of ammon-
ium chloride. The chlorine is gradually absorbed', and oily
drops of nitrogen chloride appear in the liquid. The formation
of the chloride proceeds according to the equation —

+ 4HCI

NH.Cl

3CI2

= NCI3

Ammonium

Nitrogen

chloride

trichloride

Nitrogen trichloride must be prepared in very small quantities
at a time, as it is a most violent explosive, decomposing into
its elements on the slightest provocation.

CHAPTER XXVI
BBOMINE AND IODINE

BROMINE

The chief sources of bromine are the bromides contained in
the salt deposits of Stassfurt, in Prussia. These deposits are
worked up systematically for the potassium salts which they
contain, the bromine being derived from the liquors out of
which the potassium salts have crystallised.

Bromine can be liberated from a bromide, just as chlorine
can from a chloride, by treatment with manganese dioxide
and sulphuric acid, thus —

2NaBr + MnOg + 2H2SO4 = MnS04 + NagSO^ + 2H2O + Brg

It may also be set free by passing chlorine into the bromide
solution, in the manner expressed by the following equation : —

2NaBr

CI2

= . 2NaCl

Brg

Bromide

Chlorine

Chloride

Bromine

These equations hold good for any metallic bromide, magnesium
bromide being that mostly dealt with in the actual manufacture.
If the solutions are kept near the boiling point, the bromine
distils off as vapour, which may afterwards be condensed to a
liquid.

At the ordinary temperature bromine is a dark, almost black,
liquid, which has a strong, irritating odour, resembling that
of chlorine. It is a very volatile substance, boiling at about
60**, and giving off a reddish brown vapour even at the ordinary
temperature. When shaken up with water, it only partially
dissolves, the bulk of the liquid separating out as an oily layer
on the bottom of the vessel. The solution has a colour similar
to that of bromine vapour, and is known as bromine water.
Bulk for bulk, bromine is about three times as heavy as water.

182

BROMINE AND IODINE 183

In chemical properties, bromine resembles chlorine very
closely, attacking in general those substances which are
attacked by chlorine. Thus it unites readily with most
metals, and also with non-metallic elements like sulphur or
phosphorus. It does not combine with hydrogen so readily as
chlorine does; indeed, a mixture of bromine vapour and
hydrogen may be exposed to sunlight, or brought in contact
with a lighted taper without combination occurring. The
union does take place, however, if the mixed gases are passed
through a red-hot tube. Although its power of combination
with hydrogen is thus feebler than is the case with chlorine,
bromine can yet be used as a dehydrogenising and bleach-
ing agent. Its action on alkalies is precisely the same as
that of chlorine. If the alkaline solution is cold, a liypo-
bromite is produced, if hot, a bromate.

2NaOH + Brg = NaBr + NaOBr + HgO

Bromide Hypobromite

6NaOH + 3Br2 = sNaBr + NaBrOg + sHgO

Bromide Bromate

The hypobromites are bleaching agents in the same way as
the hypochlorites. The bromates, too, resemble the chlorates
very closely. There is this difference between them, however.
The bromates on heating give no perbromate, but pass directly
into bromide and oxygen, thus —

2KBr03 = 2KBr + 3O2

Perbromic acid and perbromates are, in fact, unknown ; and
so, likewise, are oxides of bromine.

Hydrobromic acid

When sulphuric acid acts on a bromide, hydrobromic acid
is produced, just as hydrochloric acid is produced from a
chloride under similar circumstances.

NaBr + H2SO4 = NaHSO^ + HBr

A further action, however, here takes place, for the strong
sulphuric acid acts on the hydrobromic acid to some extent
as an oxidising agent, liberating bromine, thus —

2HBr 4- H2SO4 = Br2 + SO2 + 2H^0
Hydrobromic Acid Bromine

1 84

INORGANIC CHEMISTRY

No such action occurs with hydrochloric acid, as the sulphuric
acid is unable to dissolve the union between the hydrogen and
the chlorine, whilst, owing to the hydrogen and bromine being
less firmly combined, it succeeds in effecting the decomposition
of hydrobromic acid. If we take a non-oxidising acid instead
of sulphuric acid, we can liberate the hydrobromic acid from a
bromide without decomposing it. Thus, phosphoric acid and
a bromide give a phosphate and hydrobromic acid —

2NaBr + H3PO4 = Na2HP04 + 2HBr

Gaseous hydrobromic acid, however, is most conveniently

Fig. 32. — Preparation of Hydrobromic Acid.

Bromine is run from the tap-funnel A into a mixture of red
phosphorus and water contained in the flask B. The hydrobromic
acid gas which is formed is freed from bromine by a mixture of
broken glass and red phosphorus in the U tube. If it is wished to
dissolve the hydrobromic acid in water, it may be done as shown in
the figure, a Alter funnel just dipping beneath the surface of the
water being used as a delivery tub«.

prepared from bromine. When bromine is brought into
contact with phosphorus, a very vigorous action occurs, a
bromide of phosphorus being produced. This bromide of
phosphorus, if brought into contact with a little water, is at
once decomposed in accordance with the following equation : —

PBrg + 3H2O = sHBr + H3PO

Phosphorus tribroquide Hydrobromic acid Phosphorous acid

The phosphorous acid is non-volatile, and so the hydrobroiyiip

BROMINE AND IODINE 185

acid comes off alone. • In order to make hydrobromic acid,
it is not necessary to prepare phosphorus tribromide specially.
A quantity of red phosphorus is mixed with a little water, and
bromine is added drop by drop from a tap-funnel. Action at
once ensues, and hydrobromic acid comes off steadily. The
total action may be represented by the equation —

P + 3Br + 3H2O = 3HBr + H3PO

The hydrobromic acid may be freed from a little bromine
vapour which passes over with it by passing it through a tube
containing red phosphorus.

The acid thus prepared is very like hydrochloric acid. It is
a colourless gas which is extremely soluble in water, and
fumes strongly in air. It is much heavier than hydrochloric
acid, a^ its high molecular weight shows. The aqueous
solution is scarcely to be distinguished from hydrochloric
acid by its behaviour towards metallic oxides or metals, but
the two can easily be discriminated from each other by means
of chlorine. Chlorine has, of course, no action on hydrochloric
acid, but it immediately decomposes hydrobromic acid with
liberation of bromine, which can be recognised by its reddish
brown colour. The same test may be applied to all bromides,
the solutions of which, on addition of chlorine water, yield
bromine.

The bromides in other respects closely resemble the corre-
sponding chlorides. Thus silver bromide, like silver chloride,
is insoluble in water and acids ; and in general it may be said
that whatever the solubility of the chloride of a metal may be,
the solubility of the bromide of the same metal will closely
approximate to it.

IODINE

The chief sources of iodine are certain species of seaweed,
and the mother liquors derived from the crystallisation of
Chili saltpetre. The seaweed is dried, and either burned in
shallow pits, or better, distilled. The residue in either case
contains iodine in the form of iodide; The iodide is extracted
with water, and is either precipitated by passing chlorine into
the liquors, or distilled off as vapour by treating the liquor with
manganese dioxide and sulphuric acid, the actions which take

i86 INORGANIC CHEMISTRY

place being exactly analogous to those occurring in the libera-
tion of bromine from bromides.

2NaI + CI2 = 2NaCl + I2
2NaI + MnOg + 2H2SO4 = MnS04 + NagSO^ + 2H2O + Ig

In the mother liquors obtained from the crystallisation of
crude Chili saltpetre (sodium nitrate), iodine is contained in
the form of sodium iodate. This is an oxidising substance,
which yields iodine as its first reduction product (compare the
oxidation of iodine to iodic acid by means of nitric acid, p. 187).
In practice the iodine is obtained by adding a mixture of
sodium sulphite and sodium hydrogen sulphite to the iodate
liquors. These substances are oxidised to sulphate, and the
iodate is reduced to iodine, thus —

2NaI03 + 3Na2S03 + 2NaHS03 = 5Na2S04 + I2 + H2O

Iodate Sulphite Sulphate Iodine

Excess of sulphite must be avoided, for otherwise the iodine
would be further reduced to iodide. The iodine separates out
as a dark precipitate, which is pressed free from liquor and
purified by sublimation.

Iodine is a dark solid which is very sparingly soluble in
water, but freely soluble in an aqueous solution of potassium
iodide, the solution being brown in colour. When heated,
it melts at a temperature not much above 100°, and sends
off, even at that temperature, a fine violet-coloured vapour.
This vapour is very characteristic of the substance, and
when cooled condenses in shining black scales. Iodine
is usually met with in this latter form, since it is almost
invariably purified by sublimation. Iodine is much more
soluble in carbon disulphide than it is in water, so that if
the brown aqueous solution is shaken up with a little of the
disulphide, which does not itself mix with water, the iodine
leaves the water and dissolves in the disulphide, the solution
produced being of a violet colour resembling that of iodine
vapour.

Iodine acts chemically in much the same way as chlorine
and bromine, combining directly with many metals and non-
metals to form iodides. It will only combine partially with
hydrogen, however, when the two elements are brought

BROMINE AND IODINE 187

together at a high temperature. It is much more readily
oxidised than either chlorine or bromine. Thus, whilst these
substances are not attacked by strong nitric acid, iodine when
boiled with nitric acid is oxidised to iodic acid, the nitric acid
being reduced principally to nitrogen peroxide, in accordance
with the equation —

5HNO3 + I = HIO3 + 5NO2 + 2H2O

Iodine Iodic acid

The action of iodine on alkaline solutions is similar to that
of chlorine and bromine. When the alkaline solutions are
cold, iodide and hypoiodite are produced ; when the solutions
are hot, iodide and iodate. It must be remarked, however,
that the hypoiodites are extremely unstable, and pass very
soon into iodates. Iodine itself has little bleaching action,
but the hypoiodite solutions are powerful bleaching agents.

The most characteristic test for iodine is the deep blue
colour which even a trace of it will produce when brought
into contact with starch solution. The blue substance is
frequently called " iodide of starch," but there is no evidence
in support of its being a true chemical compound — it behaves
rather as if it were starch dyed with iodine. When the solution
is warmed, the colour disappears, but reappears when the
solution is cooled. This reaction is largely made use of
as an indirect test for oxidising agents (compare pp. 100, 133).
Almost any oxidising agent will liberate iodine from an iodide,
so if a moist paper impregnated with potassium iodide and
starch is brought into contact with an oxidising substance,
it speedily assumes a blue colour owing to the production of
the " iodide of starch."

Hydriodic acid

Hydriodic acid is generally prepared from iodine, phosphorus,
and water, by a reaction similar to that employed in the
preparation of hydrobomic acid. The equation is —

P + 3I + 3H2O = 3HI + H3PO3

Like hydrochloric and hydrobromic acids, hydriodic acid is a
heavy colourless gas which fumes strongly in moist air> and is

i88 INORGANIC CHEMISTRY .

extremely soluble in water. When heated to a temperature
approaching a red heat, the gas is partially decomposed into
hydrogen and iodine, thus —

2HI ^ H2 + I2

This action is a reversible one, for when the elements are
mixed at a similar temperature, partial combination results.
Hydrobromic and hydrochloric acids are not thus decomposed
by heat.

When a strong solution of hydriodic acid is exposed
to air, it rapidly darkens owing to liberation of iodine, the
oxygen of the air combining with some of the hydrogen of the
hydriodic acid to form water. Here again it is evident that
iodine parts with hydrogen much more readily than either
bromine or chlorine. Owing to this circumstance, hydriodic
acid solution is sometimes used as a reducing or hydrogen-
ising agent. Thus, if we mix a solution of iodic acid with
a solution of hydriodic acid, iodine is formed, not only by the
reduction of the iodic acid, but also by the oxidation of the
hydriodic acid, the hydrogen and oxygen of the original acids
uniting to form water.

HIO3 + 5HI = 3H2O + 3I2

Iodic acid Hydriodic acid Water Iodine

Strong sulphuric acid is also reduced by hydriodic acid, with
production of sulphurous acid and iodine —

H2SO4 + 2HI = 2H2O + SO2 + I2

For this reason strong sulphuric acid cannot be used to prepare
hydriodic acid from an iodide. It should be remembered
that dilute sulphuric acid does not act as an oxidising agent,
and in consequence it does not liberate iodine from hydriodic
acid ; indeed, the action expressed by the above equation is
reversed when much water is present — ue, sulphurous acid and
iodine produce sulphuric acid and hydriodic acid, thus —

2H2O + SO2 + I2 = H2SO4 + 2HI

The iodides bear a general resemblance to the chlorides and
bromides, but differ from them in many points. Thus, while
it may be said that if the chloride and bromide of a metal are

BROMINE AND IODINE 189

insoluble, the iodide is also insoluble, the converse statement
is not true ; for there are many insoluble iodides corresponding
to soluble bromides and chlorides. Again, where the bromides
and chlorides are colourless, the iodides are frequently coloured.
For example, mercuric chloride and mercuric bromide are
colourless and soluble, whilst mercuric iodide has a brilliant
scarlet colour and is insoluble in water.

Comparison of the Halogen Elements

Chlorine, bromine, and iodine form the natural family of
the halogens, and to them is sometimes added a fourth
element, — fluorine. Fluorine, however, diverges more from
these three elements than they do from each other, and will
not be considered here. The resemblance of the three
elements is apparent from the formulae of their corresponding
compounds, to which similar names have been given, as may
be seen in the subjoined table —

Chlorine . . CI 2 Bromine . Br 5 Iodine . . 1 2

Hydrochloric acid HCl HydrobromicacidHBr Hydriodic acid HI

Sodium chloride NaCl Sodium bromide NaBr Sodium iodide Nal

Chloric acid . HCIO3 Bromic acid . HBrOj Iodic acid . HIO,

Potassium Potassium Potassium iodate KIO3

chlorate . KClOj bromate . KBrO,

Perchloric acid HCIO4 Periodic acid HIO4

Not only are these compounds similar in their formulae —
they are also similar in their properties. The halogen elements,
therefore, may be substituted for each other without any great
change in the properties of the resulting compounds.

Bromides and iodides occur in nature as well as chlorides,
but in very much smaller quantity, so that the compounds of
bromine and iodine are comparatively rare and expensive.
Sea water, for instance, contains at least fifty times as much
chlorine in the form of chloride as it does bromine in the form
of bromide. Only mere traces of iodine exist in sea water, but
this quantity is available for the production of iodine owing to
the fact that certain seaweeds absorb and concentrate these
iodine compounds, which are afterwards found as iodides in
the ash produced when the seaweed is burnt.

When we compare the halogen elements with each other, we

I90 INORGANIC CHEMISTRY

often find a distinct gradation of properties, bromine being
usually intermediate between chlorine and iodine. Thus we
have the series of combining weights CI = 35.5, Br = 80, 1 = 1 27.
Again, at the ordinary temperature chlorine is a gas, bromine
a liquid, and iodine a solid. Comparing the depth of colour
of the vapours, we find that chlorine has least colour, and iodine
most, with bromine intermediate. With respect to chemical
activity, too, the same gradation appears. Chlorine combines
with hydrogen with great readiness, bromine less readily, and
iodine only partially at a high temperature. Conversely
hydriodic acid and iodides are easily oxidised with liberation of
iodine, hydrobromic acid and bromides require more powerful
oxidising agents, hydrochloric acid and chlorides more powerful
oxidising agents still, to effect the liberation of the halogen.
Bromine liberates iodine from iodides, but not chlorine from
chlorides, being thus intermediate in properties between the
two other halogens.

When we compare the oxygen compounds of the halogens,
however, we find that bromine no longer occupies a position
between chlorine and iodine. Thus, whilst both chlorine and
iodine have oxides, no oxides of bromine exist; and again,
although we know both perchlorates and periodates, no per-
bromates have ever been prepared.

CHAPTER XXVII

STTLPHTJB

Sulphur is an element which occurs in the uncombined state
in many volcanic districts, particularly in Sicily, from which
the bulk of our supply is derived. The native sulphur can

•''ig' 33-— Distillation cif Sulphur.
Sulphur in tfae telotl S h bailed by mcanii of a fire ia ihe
grawC. The heai fton Ibis lire Hrvei to k«p sulphur nulled
■n the reservDii /!, trom which the retorl can be replcniihed. The
sulphui vapour puses into the condeniing chamber C. and
ullimately lorms a liquid layer on its floDr.

easily be purified by distillation, since it melts at a com-
paratively low temperature, and boils below a red heat.

In the combined state sulphur chiefly occurs along with
metals in the form of metallic sulphides, or along with metals
and oxygen in the form of metallic sulphates. Some of the
sulphides — for example, iron pyrites FeSai and copper pyrites

192

INORGANIC CHEMISTRY

CuFeSg — part with a portion of their sulphur when heated
in closed vessels, commercial sulphur being frequently obtained
in this way.

The sulphur vapour produced by the distillation of crude
native sulphur is usually condensed in large brickwork
chambers. At the beginning of the distillation, the walls
of the chamber are cold, and the sulphur condenses on them
in the form of a pale yellow powder which is known as flowers
of sulphur. As the operation proceeds the walls become hot,
and the sulphur condenses not as a solid but as a liquid,
which collects on the floor of the chamber and is drawn off
into cylindrical moulds, where it solidifies on cooling to form
the ordinary roll sulphur.

When sulphur is heated to a temperature a little above that
of boiling water, it melts to an amber-coloured fluid, which

on further heating be-
^^ comes darker in colour,
and finally boils with
production of a deep
reddish brown vapour.
Different varieties of
sulphur may be ob-
tained by cooling the
heated liquid, accord-
ing to the method by
which the cooling is
effected. If the sul-
phur is cooled sud-
denly by being poured
into cold water, it forms a soft stringy material which is
known as plastic sulphur. If a quantity of it is allowed to
cool slowly in a covered vessel until a crust forms on the
surface, and if the portion which is still liquid is now poured
off through a hole broken in the crust, the interior of the
vessel will be seen to be filled with transparent brownish
yellow needles of monoclinic sulphur, which on standing
gradually lose their transparency and pass into the ordinary
lemon yellow rhombic sulphur.

Rhombic and monoclinic sulphur are crystalline; plastic
sulphur is amorphous. Both of the crystalline modifications
are soluble in carbon disulphide, whilst the amorphous

Fig. 34. — Formation of Plastic Sulphur.

Sulphur is distilled from the retort R. The vapour
condenses in the beak of the retort and flows into
cold water, the sudden chilling causing it to assume
the form of stringy masses of non-crystalline plastic
sulphur, which slowly become brittle and pass into
the ordinary form.

SULPHUR

193

Fig. 35. — Crystal of
Rhombic Sulphur.

sulphur is riot. Roll sulphur when freshly cast consists
chiefly of monoclinic sulphur, but this
variety slowly passes into rhombic sul-
phur.

Flowers of sulphur are partially crystal-
line and partially amorphous, so that on
treatment with carbon disulphide a portion
dissolves and a portion remains unaffected.
When a solution of sulphur in carbon
disulphide is allowed to evaporate, trans-
parent amber-coloured crystals of the
rhombic variety are deposited.

Sulphur, then, exists in various allotropic modifications just

as carbon does, some of these being cry-
stalline and some amorphous. The two
elements differ from each other, however,
in this respect, that whilst both the crystal-
line varieties of carbon (graphite and
diamond) can be kept for an indefinite
period, only the rhombic variety of sul-
phur is permanent at the ordinary tem-
perature.

Sulphur may also be thrown out of
solution in the form of a yellow or white
precipitate. Thus, if chlorine water is
added to a solution of sulphuretted hydro-
gen, the liquid at once becomes milky by
the precipitation of sulphur.

CL + H«S = 2HCI + S

Fig. 36. — Union of
Sulphur with Copper.

A spiral of copper wire
introduced into the va-
pour of sulphur boiling in
a test-tube becomes red-
hot from the heat de-
veloped by its combina-
tion with the sulphur to
form copper sulphide.

Sulphur precipitated from calcium penta-
sulphide by the action of an acid

CaSfi + 2HCI = CaClg + HgS + 4S

is white and very finely divided, and is
used in medicine under the name of
"milk of sulphur."

Sulphur combines readily with many
elements. For example, it burns in air
with a blue flame, forming the oxide SO 2 ; and when

194 INORGANIC CHEMISTRY

heated with metals such as iron, zinc, and copper, it
unites readily with them to form sulphides. It is attacked
by chlorine, bromine, and iodine, and at a white heat
combines with carbon to produce carbon disulphide CSg.
Many oxidising agents attack it. Thus, when heated with
concentrated nitric acid, it is converted into sulphuric acid,
and at a sufficiently high temperature it is oxidised by solid
substances such as potassium nitrate, or potassium chlorate.
On this account it is used as an ingredient of gunpowder
and of many pyrotechnic mixtures. When such mixtures
are fired, the nitrate or chlorate suddenly parts with its oxygen
to the sulphur, the action being occasionally so rapid as to
be accompanied by an explosive evolution of gas.

Oxides of Sulphur

SulplLTir dioxide, SOg. — This substance is produced when
sulphur burns in air or in oxygen. It is gaseous under
ordinary circumstances, and possesses the familiar smell of
burning sulphur, by means of which property it is most easily
recognised. The most convenient mode of preparation of
sulphur dioxide in the laboratory is by the action of a metal,
usually copper, on warm concentrated sulphuric acid, thus —

Gu + 2H2SO4 = GUSO4 + 2H2O + SO2

At the temperature of a good freezing mixture of ice and
salt, the gas condenses to a colourless liquid. The; same
liquefaction can also be brought about at the ordinary tem-
perature by the application of about two atmospheres pressure.
Corresponding to this easy condensibility, the gas is moder-
ately soluble in water, one volume of water dissolving about
40 volumes of the gas under ordinary conditions. The
aqueous solution of the gas contains sulphurous acid formed
according to the equation —

SO2 + H2O $ H2SO8

Sulphurous anhydride Sulphurous acid

Sulphurous acid, HsSOg, has not been obtained in the pure
state, since it breaks up very readily again into its anhydride
and water. It is a dibasic acid, and the salts formed by its

SULPHUR 195

neutralisation are called, sulphites. Thus the formula of
normal sodium sulphite is NagSOa. When any sulphite is
warmed with hydrochloric or sulphuric acid it is decomposed
with evolution of sulphur dioxide. For example, sodium
sulphite and hydrochloric acid react according to the
equation —

NagSOs + 2HCI = 2NaCl + HgO + SO2

Sulphur dioxide is chiefly prepared for the manufacture of sul-
phuric acid, but it also finds extensive use as a bleaching agent
and as a disinfectant.

Sulphur trioxide, SO3. — Although sulphurous acid and the
sulphites tend to take up oxygen from the atmosphere under
ordinary conditions, and become thereby converted into
sulphuric acid and sulphates —

2H2SO3 + 02 = 2H2SO4
2Na2S03 +02 = 2Na2S04,

gaseous sulphur dioxide only unites directly with oxygen under
exceptional circumstances. The two gases may be heated
together alone without union taking place, but it has been
found that the presence of various apparently inert bodies pro-
motes the union without these bodies being themselves affected.
Thus, if we pass a mixture of the gases through a tube con-
taining heated ferric oxide or finely divided platinum, combina-
tion takes place according to the equation —

2SO2 -1-02 = 2SO8

Sulphur trioxide thus produced is a gas at the temperature
of the reaction, but it readily condenses, not to a liquid, but
to fibrous silky masses, which fume strongly in air, uniting
with the moisture in it to form sulphuric acid.

Sulphuric acid

Sulphuric acid, HgSO^, is the acid which is most exten-
sively prepared on the commercial scale, most other acids
being obtained from salts by its aid. The hydrogen
which it contains is derived from water, the oxygen partly
from water and partly from the atmosphere, and the sulphur

196

INORGANIC CHEMISTRY

either from native sulphur or from a metallic sulphide
such as pyrites. The first stage in the manufacture of
sulphuric acid is the production of sulphur dioxide, which
is then made to unite with water and oxygen according to
the equation —

2S0a + O2 + 2H2O = 2H2SO4

In the most modern sulphuric acid works, sulphur trioxide
is prepared from sulphur dioxide and oxygen in the manner
indicated in the preceding section, and is then brought into
contact with water, sulphuric acid being produced.

SO3 +

Sulphur trioxide

Water

HjSO*

Sulphuric acid

By far the greatest proportion of sulphuric acid, however, is
still produced by a process which has been in use for over a
century. In this process nitric oxide NO acts as carrier of
oxygen from the air to the sulphur dioxide, instead of the pla-
tinum or ferric oxide previously mentioned. The operation
is conducted in large leaden chambers, which are supplied
with air, steam, sulphur dioxide, and small quantities of oxides
of nitrogen to make up for unavoidable loss. The actions
which go on are most simply represented by the following
equations : —

2NO

Nitric oxide

Oxygen

2NO2

Nitrogen peroxide

NO 2

Nitrogen
peroxide

SOii

Sulphur
dioxide

H2O

Water

H2SO4

Sulphuric
acid

+ NO

Nitric
oxide

The nitric oxide is oxidised by the oxygen of the air to
nitrogen peroxide, which then reacts with sulphur dioxide
and steam to form sulphuric acid, nitric oxide being at
the same time regenerated. The regenerated nitric oxide can
again take up oxygen, and start the whole process afresh.
Thus it merely plays the part of a bearer of oxygen, and
in theory a very small quantity would suffice to effect the
combination of an unlimited amount of sulphur dioxide,
oxygen, and water ; but in practice a slight loss is unavoidable,

SULPHUR 197

so that fresh quantities of nitric oxide must be supplied
from nitre.

The chamber acid produced in this way contains only
about 70% of sulphuric acid, the rest being almost entirely
water, which must be driven off by heating in leaden, glass,
or platinum vessels. The crude commercial acid prepared
from pyrites contains considerable amounts of various impuri-
ties, especially oxides of nitrogen, lead in the form of lead
sulphate, and arsenic compounds. It is frequently purified by
re-distillation.

The pure acid is nearly twice as heavy as water, bulk for
bulk, is quite colourless, and boils at a temperature over
3oo°C. On account of its comparatively high boiling point,
it can drive out more volatile acids from their salts on heating,
and is thus used in the preparation of hydrochloric, nitric, and
other acids. Its most remarkable property is the avidity with
which it absorbs water. Not only does it eagerly take up
water-vapour from the air and other moist gases, but it even
removes the elements of unformed water from many com-
pounds, thus breaking them up and destroying them. For
example, most organic compounds, which contain hydrogen
and oxygen in combination with carbon, are blackened and
charred by strong sulphuric acid. The charring consists in
the removal of the hydrogen and oxygen as water, a black
residue of carbon being left. Although no definite compound
appears to be formed, the mixing of strong sulphuric acid
with liquid water develops so much heat as often to convert
a large proportion of the water into steam.

Sulphuric acid dissolves sulphur trioxide to form what is
known as Nordhausen or faming sulphuric a^^id, which is
sometimes represented by the formula H2S2O7.

Sulphates. — Sulphuric acid is a dibasic acid and forms acid
salts, such as NaHS04, sodium hydrogen sulphate, as well as
normal salts, such as ordinary sodium sulphate, Na2S04.
Some of the most common sulphates are —

Gypsum .... CaS04,2H20

Epsom salts .
White vitriol
Green vitriol
Blue vitriol .

MgS04,7H20
ZnS04,7H20
FeS04,7H20
CuS04,5H20

Many of the sulphates are decomposed by heat, sulphur

198 INORGANIC CHEMISTRY

dioxide and oxygen being evolved, and metallic oxide left
behind. Thus copper sulphate decomposes according to the
following equation : —

2CUSO4 = 2CuO + 2SO2 + O2

Copper sulphate Copper oxide Sulphur dioxide Oxygen

Ferrous sulphate or green vitriol when heated yields, amongst
other products, sulphuric acid, which was first prepared in this
way, and from its source and oily appearance received the
name of oil of vitrioL

SulplLuric acid as an ozidismg agent. — When sulphuric
acid is raised to a red heat, as may be done for example by
throwing the acid on red-hot bricks, it is decomposed in much
the same way as copper sulphate, thus —

2H2SO4 = 2H2O + 2SO2 + O2

Sulphuric acid Water Sulphur dioxide Oxygen

If a substance capable of uniting with the oxygen is also
present, the decomposition is effected at a much lower temper-
ature, so that concentrated sulphuric acid frequently behaves
as an oxidising agent, being reduced in the action to sulphurous
acid, or to sulphur dioxide and water. Thus, if pure sulphuric
acid is warmed with carbon, sulphur, or a metal such as zinc or
copper, it oxidises these substances according to the following
equations : —

C + 2H2SO4 = CO2 + 2SO2 + 2H2O

Carbon Carbon dioxide

S + 2H2SO4 = SO2 + 2SO2 4- 2H2O

Sulphilr Sulphur dioxide

Zn + 2H2SO4 = ZnS04 + SO2 + 2H2O

Zinc Zinc sulphate

Cu + 2H2SO4 = CUSO4 + SO2 + 2H2O

Copper Copper sulphate

As has already been stated, the action of strong sulphuric
acid on copper is used in the laboratory for the preparation of
sulphur dioxide.

It is only if the sulphuric acid is pure or mixed with a
small quantity of water that it acts in this way as an oxidising

SULPHUR 199

agent; dilute sulphuric acid has no oxidising properties.
Thus, dilute sulphuric acid leaves carbon, sulphur, and copper
quite unaffected ; and although it does act on zinc, no sulphur
dioxide is produced, hydrogen being evolved in its stead.

Zn + H2SO4 = ZnS04 + H

In dilute solution the tendency is rather for sulphurous acid
and sulphites to pass into sulphuric acid and sulphates (p. 195).
A solution of sodium sulphite, for example, absorbs oxygen
from the air, and is gradually converted into sodium sulphate —

2Na2S03 + 02 = 2Na2S04

Sodium sulphite Sodium sulphate

Similarly, sulphurous acid can often be employed as a reducing
agent (p. 100).

A good example of the effect of water in determining the
relative stability of the two oxygen acids of sulphur is afforded
by the reversible action (p. 188) : —

H2SO4 + 2HI 5> 2H2O + SO2 + I2

Sulphuric acid Hydriodic acid Water Sulphur dioxide Iodine

If little or no water is present, the sulphuric acid acts as an
oxidising or dehydrogenising agent, removing hydrogen from
the hydriodic acid and being itself converted into sulphur
dioxide. If much water is present, this reaction does npt
occur ; on the contrary, sulphur dioxide and iodine regenerate
sulphuric acid and hydriodic acid.

Thiosulpliates

Just as sodium sulphite unites with oxygen to form sodium
sulphate, so also does it unite with sulphur when a solution of
it is boiled with that substance. The product of the action is
called sodium thiosulphate, and its formation may be re-
presented by the equation —

Na2S03 + S = Na2S203

Sodium sulphite Sulphur Sodium thiosulphate

The prefix thio is used to indicate substitution of sulphur for
oxygen, and a reference to the formulae will show that sodium

aoo INORGANIC CHEMISTRY

thiosulphate is related to sodium sulphate by the replacement of
one-fourth of its oxygen by an equivalent amount of sulphur.

The thiosulphuric acid from which the thiosulphates are
derived is a dibasic acid, which, however, cannot be obtained
in the pure state, owing to its tendency to decompose accord-
ing to the equation —

H2S2O3 == H2SO3 + S

Thiosulphuric acid Sulphurous acid Sulphur

Thus, if hydrochloric acid is added to a solution of sodium
thiosulphate, the solution speedily becomes milky owing to the
separation of sulphur, and the smell of sulphur dioxide is
perceptible, these substances being formed by the decomposi-
tion of the thiosulphuric acid originally liberated.

Sodium thiosulpliate is largely used in photography under
the name of "hypo," derived from the name hyposulphite of
soda, which was formerly applied to this compound. Its
employment in photography depends on the property that its
solutions readily dissolve many silver salts which are insoluble
in water. Thus silver bromide dissolves in hypo solution
with production of sodium bromide and sodium silver thio-
sulphate, which are both soluble in water. Note that this is
an exception to the rule given on p. 69.

AgBr + NagSgOs = NaBr -h NaAgSgOg

Silver Sodium thiosulphate Sodium Sodium silver

bromide bromide thiosulphate

Sulphuretted hydrogen

When iron filings and flowers of sulphur are heated together,
they unite to form a sulphide of iron, which is different from
the iron pyrites found as a mineral. It contains only half as
much sulphur as pyrites, and is known as ferrous sulphide,
its formula being FeS. This substance is the usual source
of sulphuretted hydrogen in the laboratory, for when treated
with dilute hydrochloric or sulphuric acid, it decomposes as
follows : —

FeS -H H2SO4 = H2S + FeSO^

Ferrous Sulphuric Hydrc^en Ferrous

sulphide acid sulphide sulphate

The hydrogen sulphide, or sulphuretted hydrogen, thus pre-

SULPHUR 20I

pared, is not quite pure, being usually mixed with a little
hydrogen derived from a small quantity of metallic iron which
the ferrous sulphide contains, but this impurity as a rule is of
no moment.

Hydrogen sulphide is easily recognisable by its unpleasant
odour, which resembles that of rotten eggs. It is a colourless
gas, somewhat heavier than air, is moderately soluble in water
(p. 112), and is condensable to a liquid at the ordinary tempera-
ture. In a plentiful supply of air it burns with formation of
water and sulphur dioxide : if the supply of air is defective, the
hydrogen burns in preference to the sulphur. The following
equations represent the complete and partial combustion of
hydrogen sulphide : —

2H2S + 3O2 = 2H2O + 2SO2
2H2S +02 = 2H2O + 2S

A solution of hydrogen sulphide on exposure to the air
rapidly becomes turbid, owing to the deposition of sulphur
caused by partial oxidation, the equation for the action being
the last one of the preceding pair.

Sulphuretted hydrogen is a weak acid, like carbonic acid,
which it also resembles in being dibasic. The normal sulphide
of sodium has the formula Na2S, but in addition to this, there
is a sulphide NaHS, which is called sodium hydrogen sulphide,
or sodium hydrosulphide. The solutions of these sulphides,
like those of the corresponding carbonates, have an alkaline
reaction, due to their partial decomposition by water, with
liberation of sodium hydroxide (p. 141).

All the metallic sulphides, except those of the alkali metals,
are insoluble in water. Some, however, are decomposed by
water with formation of soluble products. Thus, calcium
sulphide on being brought into contact with water slowly
splits up according to the following equation : —

2CaS +

2HjO =

= Ca(HO)j

+ Ca(HS)j

Calcium

Water

Calcium

sulphide

hydroxide

hydrosulphide

the hydrosulphide being easily soluble, and the hydroxide
sparingly soluble in water. On account of the insolubility
of the majority of metallic sulphides, and on account of the
characteristic colours which many of them possess, sulphuretted

202 INORGANIC CHEMISTRY

hydrogen is a valuable reagent in the laboratory for identifying
the metallic radicals in salts. For example, if we add
sulphuretted hydrogen to solutions of the sulphates of
copper, zinc, and cadmium, we obtain precipitates of the
corresponding sulphides, in accordance with the equations —

CUS04

H2S

CuS

H2SO4

ZnS04

H2S

ZnS

H2SO4

CdS04

H2S

CdS

H2SO4

Of these, copper sulphide is black, zinc sulphide white, and
cadmium sulphide yellow. The different solubility of the
sulphides in dilute acid, too, affords a valuable means of
separating the metallic radicals into well-defined groups.
Thus, zinc sulphide dissolves in very dilute hydrochloric
acid, whilst copper and cadmium sulphides are practically
unaffected by an acid of the same concentration. It must
be borne in mind that, in such a case, the sulphide does
not dissolve as such in the acid, but is decomposed by it,
with formation of a soluble salt and sulphuretted hydrogen —

ZnS + 2HCI = ZnCla + HgS

Zinc sulphide Zinc chloride

The presence of a very small amount of hydrochloric acid will
thus prevent the precipitation of zinc sulphide from a zinc
salt by means of sulphuretted hydrogen, but will not interfere
with the precipitation of copper or cadmium sulphides under
the same conditions, thus affording a method of separating
zinc from copper and cadmium.

When heated, or roasted^ in air, many of the metallic
sulphides are converted into sulphates, though this action
is also usually accompanied by the formation of oxides.
For example, lead sulphide on being roasted at a high
temperature is oxidised as follows : —

PbS + 2O2 = PbS04

Lead sulphide Lead sulphate

2PbS + 3O2 = 2PbO + 2SO2

Lead sulphide Lead oxide Sulphur dioxide

Both these actions are made use of in the extraction of
lead from the sulphide, which is its chief ore (p. 230).
On the other hand, the metallic sulphates can usually be

SULPHUR 203

converted into the corresponding sulphides by heating them
with carbon at a red heat, thus —

BaS04 + 4C = BaS + 4CO

Barium sulphate Barium sulphide

Chlorides of Sulphur

The ordinary chloride of sulphur, which is largely used in
vulcanising rubber, has the formula S2CI2J and is usually
called the monochloride, to distinguish it from the others,
which have proportionately more chlorine. It is prepared
by the direct union of the elements, chlorine being passed
into a vessel containing gently heated sulphur. The chloride
is formed according to the equation —

2S + CI2 = S2CI2

and distils over at the temperature of the operation.

It is a yellow liquid with a peculiar unpleasant odour, and
boils at a temperature somewhat above the boiling point of
water. It sinks in water and is slowly decomposed by it,
forming hydrochloric acid, sulphurous acid, and sulphur, thus —

2S2CI2 + 3H2O = 4HCI + H2SO3 + 3S

The other chlorides, SCI 2 and SCI 4, are unstable and of
no practical utility.

Carbon disulphide

Carbon disulphide (or carbon ^/sulphide, as it is still usually
called) differs from the metallic sulphides as much as carbon
dioxide differs from the metaUic oxides. It is formed by the
direct union of sulphur vapour and carbon (in the form of
charcoal or coke) at a bright heat, thus —

C + 2S = CS2

Carbon Sulphur Carbon disulphide

Carbon disulphide is a volatile liquid which will not mix
with water, and on which water floats, on account of the
higher specific gravity of the disulphide. When perfectly
pure it has a pleasant smell resembling that of chloroform,
but on standing it rapidly acquires a very offensive odour,
by means of which it is easily recognised. It is chiefly useful
as a solvent for some substances which do not dissolve in
water — e,g, fats, oils, sulphur, phosphorus.

204 INORGANIC CHEMISTRY

It burns readily in oxygen or air with a blue flame, the
carbon becoming carbon dioxide, and the sulphur, sulphur
dioxide —

CSa + 3O2 = CO2 + 2SO2

In a deficient supply of oxygen, the carbon burns in preference
to the sulphur.

Comparison of Sulphur and Oxygen

Although widely divergent in their physical properties,
oxygen and sulphur exhibit so many points of analogy in
their compounds that chemists are in the habit of classifying
them along with each other in the same natural family of
elements. It must be admitted, however, that the resemblance
between the corresponding compounds is mostly a resemblance
in formulae and not in properties, differing therefore in char-
acter from the resemblance between the halogen elements.
From the following table it is evident that sulphur is capable
of taking the place of oxygen, but the compounds thus
produced are often very dissimilar in physical and chemical
characters : —

Oxygen compounds Sulphur compounds

HgO . . Water HgS . . Sulphuretted hydrc^en

NaOH . . Sodium hydroxide NaSH . Sodium hydrosulphide

CaO . . . Calcium oxide CaS . . Calcium sulphide

NajSO^. . Sodium sulphate NagSgOa Sodium thiosulphate

CO 3 . . . Carbon dioxide CS^ . . Carbon disulphide

Whilst water is a neutral odourless liquid, sulphuretted hydrogen
is an acid offensive gas. Sodium hydroxide is a very powerful
alkali: sodium hydrosulphide has scarcely any alkaline properties.
Carbon dioxide is a gas under ordinary conditions : carbon
disulphide is a heavy liquid, and so on. We have here, then,
a formal resemblance between oxygen and sulphur, but little
real resemblance in properties. It is important that the
student should note this point, for much of chemical classifica-
tion is based on resemblances which are more formal than real.

The following scheme indicates by means of arrows how the
various compounds of sulphur are usually derived from each
other, and the student is strongly recommended to draw up for
himself similar tablesingreaterdetailforthis and otherelements: —

SULPHUR

205

•»-»

^^-^

•

-o

PJH

•^

•rH

^■^

<!-•

CO
<L)

CHAPTER XXVIII

PHOSPHOBUS

Phosphorus always occurs in nature in the oxidised state
as phosphate, and practically the only source of phosphorus
compounds is calcium phosphate Ca3(P04)2. This substance
is found nearly pure in certain minerals, and can be prepared
in quantity from bones, which in the dry state contain fully
half their weight of calcium phosphate, and when burned in air
leave behind bone-ash consisting chiefly of this compound.

In order to prepare phosphorus, calcium phosphate is heated
in an electric furnace together with silica (in the form of sand)
and charcoal. At the high temperature of the furnace, the
carbon reduces the phosphate to phosphorus, and the silica,
being an acid anhydride, unites with the calcium oxide to
form calcium silicate. The action may be represented thus —

Ca3(P04)2 + sSiOa + 5C = 2P + sCaSiOs + 5CO
Calcium Silica Carbon Phos- Calcium

phosphate phorus silicate

Or calcium phosphate may first be converted into phosphoric
acid by the action of sulphuric acid. The metaphosphoric acid,
on heating with charcoal alone in retorts contained in an ordi-
nary furnace, is reduced by the carbon, with production of
phosphorus, carbon monoxide, and hydrogen.

2HPO3 + 6C = 6C0 + H2 + 2P

The phosphorus in either case distils over in the form of
vapour, and is condensed in water. It is then melted under
warm water and cast into sticks. When quite pure, phos-
phorus is perfectly colourless and transparent, but as actually
obtained it generally possesses a yellow colour, and is known
as yellow phosphorus. At the ordinary temperature yellow
phosphorus is a soft waxy substance which can be scratched by
the nail or cut with a knife. It melts at 44°, and if exposed
to air at that temperature, invariably takes fire. The heat

206

PHOSPHORUS 207

developed by friction in cutting it at the ordinary temperature
is often sufficient to inflame it, so that it should always be cut
under water. It is customary indeed to keep phosphorus
permanently under water, so as to avoid the liability to accident
from its easy inflammability.

Yellow phosphorus has a characteristic smell — the smell of
ordinary matches — ^and is seen to glow in the dark when ex-
posed to moist air. It is soluble in carbon disulphide, and is
deposited in the crystalline state when the solvent is allowed
to evaporate slowly.

Yellow phosphorus even at the ordinary temperature slowly
combines with the oxygen of the air, and may be used for
removing oxygen from a mixture of gases (compare p. 124).
When the phosphorus exposes a large surface to the atmos-
phere, the oxidation may proceed so rapidly as to raise the
temperature to the ignition point of phosphorus, which then
bursts into flame. Such a large surface may be secured by
allowing a little of the solution in carbon disulphide to
evaporate on a piece of filter paper, a small quantity of
phosphorus being thus spread over a great space.

Ordinary matches in this country are usually tipped with
a mixture of phosphorus, potassium chlorate, and glue. The
end of the splint is first of all dipped in melted paraffin wax,
and then into a paste made of the above substances together
with a small quantity of fine sand, and vermilion as a colour-
ing matter. When the match is rubbed on a rough surface,
the friction, which is increased by the presence of the sand,
is sufficient to raise the temperature of the composition to a
point at which the phosphorus takes fire, most of the oxygen
being supplied by the potassium chlorate. The flame of
phosphorus under these conditions is not capable of igniting
wood directly. It will, however, ignite paraffin, which in its
turn is able to ignite the wood. The glue is present merely
to hold the composition together, and fix it to the splint,
playing no essential part in the chemical action.

There is another variety of phosphorus, red phosphorus,
which differs greatly in its properties from yellow phosphorus.
Although this variety is sometimes called amorphous phos-.
phorus, it is in reality crystalline. It may be prepared by
heating phosphorus to a temperature of about 240°. If
a very small quantity of iodine is added to the liquid phos-

2o8 INORGANIC CHEMISTRY

phorus, the conversion occurs at a temperature of about

200\

Red phosphorus differs from yellow phosphorus, not only
in its crystalline form and physical properties, but also greatly
in chemical activity. It has no smell and no poisonous
action, it does not glow in the dark, it is insoluble in carbon
disulphide, and does not ignite in the air until warmed to a
temperature exceeding 200".

On account of its non-poisonous properties and smaller
liability to ignition, red phosphorus is used in the manufacture
of safety matches, or rather of the surface on which the
safety matches are struck. The tip of the safety match re-
sembles that of an ordinary match, with the exception that
the combustible substance in it is not phosphorus, but anti-
mony sulphide, SbgSa. When drawn along the striking
surface, which contains amorphous phosphorus, a little of the
phosphorus ignites at the point of contact, where it comes
into contact with the potassium chlorate in the head of the
match. The combustion, however, is not transmitted to the
rest of the amorphous phosphorus on the prepared surface,
but to the mixture on the match head, which contains both
the combustible sulphide and the chlorate to supply the
necessary oxygen.

Red phosphorus is thus much less active chemically than
yellow phosphorus, yet the difference is only one of degree.
The two substances enter into exactly the same combinations,
but the yellow phosphorus does so with greater readiness, and
at a lower temperature.

Besides combining with oxygen and the halogens, phos-
phorus combines with some metals to form />Aosp/udeSy which
are in many respects analogous to the sulphides.

Phosphorus in the state of vapour has a density correspond-
ing to the molecular formula P4.

Oxides of Phosphorus

The two chief oxides of phosphorus are the trioxide and the
pentoxide. When phosphorus bums in the air or in oxygen,
dense white fumes are produced, which consist of a mixture of

PHOSPHORUS 209

these two oxides, the latter predominating. Their formation
is represented by the equations —

30.

2P208

Phosphorus trioxide

502

2P2O5

Phosphorus pentoxide.

Phosphorus trioxide, PiOg. — In order to obtain this oxide,
the phosphorus is burned in a defective supply of air, and the
fumes passed through a glass tube containing a plug of glass
wool. This serves to stop the solid pentoxide, but permits the
trioxide to pass on into a condensing vessel. The trioxide
is liquid on a warm day, solid on a cold day. Its vapour
has a density corresponding to the molecular formula P40fl,
although its name is derived from the simpler formula P2O3.
It unites slowly with cold water to form phosphorous acid, for
which reason it is sometimes called phosphorous anhydride.
The equation representing this action is —

P40fl + 6H2O = 4H3PO8

Phosphorus trioxide Phosphorous acid.

When warmed in oxygen, the trioxide burns to form pentoxide —
P^Ofl + 2O2 « 2P2O6

Trioxide Pentoxide

Phosphorus pentoxide, P2O6. — When the supply of air or
oxygen in which phosphorus burns is plentiful, the pentoxide
is produced. This substance is usually met with as a white
powder, which possesses a great attraction for moisture. If
exposed to the air it deliquesces to form a syrupy mass, and if
thrown into water it dissolves with a hissing noise, owing tp
the heat produced by its combination with the water. The
action which takes place is represented by the following
equation : —

P2O5 + H2O = 2HPO3

Phosphorus pentoxide Metaphosphoric acid.

Since the substance produced by its union with water is a
variety of phosphoric acid, phosphorus pentoxide is frequently
C9X\edphosphoric anhydride. Phosphoric anhydride is largely em-
ployed in the laboratory for drying gases, etc., when it is essential
o

iio INORGANIC CHEMISTRY

to get rid of the last traces of moisture. Not only, however,
will it remove water actually present in a mixture — it will, like
sulphuric acid, remove the elements of water from a compound ;
it will even remove the elements of water from sulphuric acid
itself, thus —

HjjSO*

+ PaOs

SO3 + 2HPO3

Sulphuric

Phosphoric

Sulphuric Metaphosphoric

acid

anhydride

anhydride acid

Phosphoric anhydride is, in fact, one of the most powerful
dehydrating agents with which we are acquainted.

Phosphorus and the Halogens

Yellow phosphorus when brought into contact with the
halogens at the ordinary temperature, unites with them
spontaneously, evolving heat, and forming a phosphorus
chloride, bromide, or iodide. Red phosphorus is also readily
attacked by bromine and chlorine at the ordinary temperature,
but requires to be slightly warmed before it unites with iodine.

Phosphorus trichloride, PCI3. — This substance, formed
according to the equation —

2P + 3CI2 = 2PCI3

is a fuming liquid, which boils at a temperature lower than the
boiling point of water. When poured into water it is rapidly
decomposed with formation of hydrochloric acid and phos-
phorous acid —

PCI3 + 3H2O = 3HCI + H3PO3

Phosphorus trichloride Phosphorous acid.

Phosphorus pentachloride, POlg. — When excess of chlorine
is used, or when the trichloride is exposed to the action of
chlorine, the pentachloride of phosphorus is produced —

2P + 5CI2 = 2PCI6

PCI3 + CI2 <> PCI5

Unlike the trichloride, the pentachloride is a yellow crystalline
solid. When converted into vapour the pentachloride dis-
sociates into trichloride and chlorine, which recombine when
the vapour is cooled. The action expressed by the second of
the above pair of equations is thus reversible.

PHOSPHORUS 211

The action of water on the pentachloride is vigorous, and
similar in character to the action on the trichloride, ordinary
phosphoric acid being formed instead of phosphorous acid —

PCI5 + 4H2O = 5HCI + H3PO4
Phosphorus pentachloride Phosphoric acid.

If only a small quantity of water is used in the decomposition,
phosphortLS oxychloride, a Uquid substance containing both
oxygen and chlorine, is produced, thus —

PCI5 + H2O = 2HCI + POCI3

Phosphorus pentachloride Phosphorus oxychloride.

This oxychloride on further treatment with water is converted
into phosphoric acid.

Bromides of phosphorus. — There are two bromides of phos-
phorus, the tribromide, PBrg, and the pentabromide, PBrg,
which resemble the chlorides very closely, both in physical
and chemical properties. When acted upon by water they pro-
duce phosphorous and phosphoric acids respectively, together
with hydrobromic acid, which here appears in place of hydro-
chloric acid —

+ H3PO3

Phosphorous acid

+ H3P0,

Phosphoric acid

These actions are used in the preparation of gaseous hydro-
bromic acid.

Iodides of phosphorus. — The chief iodide of phosphorus has
the formula PI2 or P2I4. It is a reddish solid which is decom-
posed by water with formation of red phosphorus, phosphorous
acid,, and hydriodic acid, for the preparation of which it is
mostly employed (p. 187).

Oxygen Acids of Phosphorus

Phosphorus forms a series of oxygen acids, the chief of which
are noted below, together with their formulae and the names of
their salts —

PBr., +

3H,0 =

= 3HBr

Tribromide

PBrg +

4H2O =

= sHBr

Pentabromide

Acid

Salt

Phosphoric acid

H3PO4

Phosphate

Phosphorous acid .

H3PO3

Phosphite

Hypophosphorous acid .

HsPOa

Hypophosphite

212 INORGANIC CHEMISTRY

By far the most important of these is the most highly oxidised
— namely, phosphoric acid.

Phosphoric acid. — An impure phosphoric acid is prepared
from the natural calcium phosphate, or from bone ash, by
decomposing these substances with sulphuric acid, in accord-
ance with the equation —

Ca3(P04)2 + 3H2SO4 = 2H3PO4 + 3CaS04

Calcium phosphate Phosphoric acid Calcium sulphate

The calcium sulphate is insoluble and may be separated from
the phosphoric acid, which remains in solution. To obtain
pure phosphoric acid, phosphorus is oxidised by boiling with
nitric acid, thus —

3P + 5HNO3 + 2H2O = 3H3PO4 + 5NO

Phosphorus Nitric acid Phosphoric acid Nitric oxide

The excess of nitric acid is driven off by evaporation, the
heating being usually continued until the oithophosphoric acid
(ordinary or tribasic phosphoric acid) is converted into meta.-
phosphoric acid (glacial phosphoric acid) by loss of water.

H3PO4 = HPO3 + H2O

Orthophosphoric acid Metaphosphoric acid

The glacial phosphoric acid is a glassy mass which is often cast
into the form of sticks. If the metaphosphoric acid is now
dissolved in water and the solution boiled, it is reconverted
into orthophosphoric acid —

HPO3 + H2O = H3PO4

Metaphosphoric acid Orthophosphoric acid

which may be crystallised out of the solution.

The solution obtained by dissolving phosphorus pentoxide
in water is a solution of metaphosphoric acid, which may
similarly be converted into orthophosphoric acid by boiling.

If orthophosphoric acid is gently heated, an acid can be
obtained from it which is different from the original acid and
also from metaphosphoric acid. This acid is called pyrophos-
phoric acid, and its formation may be represented by the
equation —

2H3PO4 = H4P2O7 + H2O

Orthophosphoric acid Pyrophosphoric acid

PHOSPHORUS 213

Pyrophosphoric acid on further heating loses water and is
converted into metaphosphoric acid —

H4P2O7 = 2HPO8 + H2O

Pyrophosphoric acid Metaphosphoric acid

Conversely, pyrophosphoric acid is formed as an intermediate
product when orthophosphoric acid is produced from metaphos-
phoric acid by heating with water. The equations are the
reverse of those given above — namely,

2HPO8 + H2O = H4P2O7
H4P2O7 + H2O = 2H3PO4

Pyrophosphoric acid is thus exactly intermediate between
orthophosphoric acid and metaphosphoric acid, being con-
verted into the former by the addition of one formula weight
of water, and into the latter by the removal of one formula
weight of water.

The relations of these acids are perhaps rendered most
clearly evident if we consider them as consisting of water and
phosphoric anhydride in different proportions. Thus we have

P205 .

• • . .

Phosphoric anhydride

PjOg, H^O

or 2HPO3 .

Metaphosphoric acid

P^Og, 2HjO

or H4P267 .

Pyrophosphoric acid

PjOg, 3H2O

or 2H3PO4 .

Orthophosphoric acid

There are two points to be noted in connection with this
mode of viewing the phosphoric acids. If phosphoric anhydride,
metaphosphoric, or pyrophosphoric acids are left in contact
with water for a sufficient length of time, they will eventually
be converted into orthophosphoric acid — ue, into that form
which has the maximum amount of water in its composition.
On the other hand, if water is driven off from any of the
phosphoric acids by heating, the ultimate product is metaphos-
phoric acid, and not phosphoric anhydride as we might expect.
It is impossible by heating to procure phosphoric anhydride
from any of the phosphoric acids.

Each of the phosphoric acids has its own salts, but, like
the corresponding acids themselves, the meta and pyro-
phosphates pass into orthophosphates when left for a long
time in contact with water, or, more rapidly, when boiled with
water.

214 INORGANIC CHEMISTRY

Orthophosphoric acid is a tribasic cicid, and therefore forms
three sets of salts — namely, normal salts and two sets of acid
salts. We are thus acquainted with three orthophosphates of
sodium — normal sodium orthophosphate, Na3P04 ; disodium
hydrogen orthophosphate, Na2HP04 ; and sodium dihydrogen
orthophosphate, NaH2P04. The normal salt yields a strongly
alkaline solution, the second or mon-acid salt yields a feebly
alkaline solution, and the third or di-acid salt yields an acid
solution. The common phosphate of soda is the mon-acid
salt, Na2HP04.

The action of heat on the various orthophosphates is of
interest. If the base or basic oxide from which the phosphates
are derived is capable of resisting heat, as sodium hydroxide
is, for instance, the following rules hold good. The normal
phosphate is unaffected by heat, the mon-acid phosphate is
converted into a pyrophosphate, and the di-acid phosphate is
converted into a metaphosphate. The equations for the
sodium salts are as follows : —

2Na2HP04 = Na4P207 + HgO

Mon-acid salt Pyrophosphate

NaH2P04 = NaPOs + H2O

Di-acid salt Metaphosphate

All the normal orthophosphates, except those of the alkalies,
are practically insoluble in water. Bone ash and natural
calcium phosphate are extensively used as manure to supply
phosphorus to plants, without which they do not thrive. On
account of its insolubility, the calcium phosphate in this form
acts but slowly, so that " soluble phosphate " is often employed
in its stead. This soluble phosphate or superphosphate is the
di-acid calcium phosphate which, like most acid salts, is soluble
in water. It is prepared from the normal phosphate by
treating it with sulphuric acid, the quantities being chosen in
accordance with the equation —

Ca3(P04)2 + 2H2SO4 = CaH4(P04)2 + aCaSO^

Normal phosphate Soluble phosphate

The calcium sulphate being insoluble can easily be separated
from the solution of superphosphate if required.

The orthophosphates can readily be distinguished from the

PHOSPHORUS 215

meta and pyrophosphates by means of the colour of the silver
salts. Solutions of orthophosphates yield, with silver nitrate,
a bright yellow precipitate of normal silver orthophosphate :
meta and pyrophosphates, under the same conditions, give
white precipitates.

Phosphorous acid, H^POg. — This acid may be prepared by
dissolving the oxide P40e in water, or by the action of water
on the trichloride —

PCI3 + 3H2O - H.3PO3 + 3HCI

Trichloride Phosphorous acid

The hydrochloric acid may be driven off by heat, the phos-
phorous acid being obtained in the crystalline form when the
solution cools. Phosphorous acid when heated undergoes
decomposition. A part of it is oxidised to phosphoric acid at
the expense of another part which is reduced to phosphine,
thus —

4H.3PO3 . ^ 3H3PO4 + PH3
Phosphorous acid Piiosphoric acid Phosphine

In solution, phosphorous acid acts as a reducing agent, taking
up an atom of oxygen, and being converted into phosphoric
acid.

The phosphites derived from phosphorous acid are also
reducing agents.

Hypophosphoroos acid, H3PO2. — When phosphorus is boiled
with an alkali, it is partially oxidised to a salt of hypophos-
phorous acid, and partially hydrogenised to phosphine. Thus
the action of a boiling solution of caustic soda on phosphorus
is represented by the equation —

3NaOH + 4P + 3H2O » 3NaH2P02 + PH3

Caustic soda Phosphorus Sodium hyphosphite Phosphine

A solution of barium hypophosphite may be formed in the
same way by boiling phosphorus with barium hydroxide solu-
tion. If the requisite quantity of sulphuric acid is added,
insoluble barium sulphate is produced, from which the solution
of hypophosphorous acid may be separated by filtration —

Ba(H2P02)2 + H2SO4 = 2H3PO2 + BaSO^

Barium hypophosphite Hypophosphorous acid

2l6

INORGANIC CHEMISTRY

Hypophosphorous acid decomposes like phosphorous acid
when heated, forming phosphoric acid and phosphine —

2H3PO2 - HsPO^ + PH3

Hypophosphorous acid Phosphoric acid Phosphine

It also resembles phosphorous acid in being a powerful reducing
agent, taking up oxygen to become phosphoric acid.

Although it has as much hydrogen as phosphoric acid in
its formula, only one-third is replaceable by a metal — /.^. it is
a monobasic acid, and forms only one series of salts. Sodium
hypophosphite, Na*H2P02', is used in medicine.

Phosphine, PH3

We have seen that when phosphorus is boiled with an alkaline

solution, and when phosphorous or hypo-
phosphorous acid is heated, the substance
phosphine PH3 is produced. This sub-
stance, which is often called phosphuretted
hydrogen^ resembles, in some respects,
sulphuretted hydrogen. Thus when passed
through solutions of many metallic salts,

it forms precipitates of phosphides analo-

Fig. 37.— Formation S°"s to the sulphides ; and just as sul-
of Smoke Rings phuretted hydrogen may be prepared by

decomposing sulphides by means of acids,
so phosphuretted hydrogen may be made

from Phosphine.

Some calcium phos-
phide is placed at the u ~j " • t" i*^* j • ' • m

bottom of a conical glass by Qccomposmg phosphidcs m a Similar

As'^\*hL\Se'^*o^f'*'lf^^^ manner. Some phosphides, like some

phine rise to the surface carbidcs, are cvcn dccomposablc by

rion" '?f"%hiS '^mo'k.' ^ater. For example, calcium phosphide

(oxides of phosphorus) whcn thrown into water decomposes with

fbrm of* vortex ringVas* cvolution of phosphinc (comparc the

sumed by the escaping action of watcr On calcium carbide,

bubbles of gas. v '

p- 152)—

CasPj + 6H2O = 3Ca(OH)2

Calcium phosphide

Phosphine is a gas which has an unpleasant odour, and is
highly poisonous. It burns readily in air or oxygen with a
brilliant flame, dense fumes of phosphorus pentoxide or of
phosphoric acid being produced,

-H 2PH3
Phosphine

PHOSPHORUS 217

When phosphine is prepared by any of the methods men-
tioned above, it is liable to contain small quantities of another
compound of phosphorus and hydrogen — namely, P2H4,
which is usually called liquid phosphuretted hydrogen. Now
this compound is spontaneously inflammable — />. takes fire
when brought into contact with air or oxygen, even at the
ordinary temperature. It therefore inflames the gaseous phos-
phuretted hydrogen with which it is mixed, although this
substance is not itself inflammable at the ordinary tempera-
ture. This may be proved by passing the gas prepared by
any of the above methods through a layer of turpentine, which
dissolves the vapour of the liquid phosphuretted hydrogen,
and allows the gaseous phosphine to pass on. Before being
washed with the turpentine the gas is spontaneously inflam-
mable ; after washing it is so no longer.

When the unpurified gas is made to bubble slowly through
water, each bubble, when it rises to the surface and comes in
contact with the air, takes fire, producing white fumes in the
form of a vortex ring, as is indicated in the figure.

Phosphine unites with gaseous hydriodic acid to form a
crystalline solid, called phosphonium iodide —

PH3 -h HI = PH4I

Phosphine Phosphonium iodide

This action may be compared with the union of ammonia and
hydriodic acid to form ammonium iodide —

NH3 -h HI = NHJ

Ammonia Ammonium iodide

Phosphonium iodide, like ammonium iodide, is a true salt,
so that phosphine acts in this respect as an anhydrous base,
like ammonia. Just as ammonia may be liberated from
ammonium iodide by treatment with caustic soda, so pure
phosphine may be liberated from phosphonium iodide in like
manner —

PHJ + NaOH = NaT -h HgO + PH3
NH4I + NaOH = Nal + HgO + NH3

Although phosphine resembles ammonia in this respect, the
analogy stops here, Phosphine is scarcely soluble in water

2i8 INORGANIC CHEMISTRY

and does not turn red litmus blue; ammonia, on the other
hand, is excessively soluble, yielding a strongly alkaline solu-
tion. Phosphine is easily inflammable : ammonia will not
burn in air, unless heat is constantly supplied to enable the
action to take place.

Comparison of Phosphorus with Nitrogen and Sulphur

Nitrogen and phosphorus are classed together in the same
family of elements. The resemblance between them is,
however, by no means close, and is generally confined to a
practically formal similarity in the case of a few compounds.
We have, in fact, here much the same kind of relation as we
found between oxygen and sulphur. The following table
shows the formal resemblance which exists : —

NH3 Ammonia PH3 Phosphine

NH4I Ammonium iodide PH4I Phosphonium iodide

NCI 3 Nitrogen trichloride PCI3 Phosphorus trichloride

NgOg Nitrogen pentoxide P2O5 Phosphorus pentoxide

HNO3 Nitric acid HPO3 Metaphosphoric acid

The actual differences between ammonia and phosphine have
already been insisted on. Nitrogen trichloride is not
decomposed by water, and is one of the most explosive
compounds with which we are acquainted. Phosphorus
trichloride, on the other hand, has no tendency to explode,
and is at once decomposed by water. Nitrogen pentoxide
decomposes when heated into the peroxide and oxygen :
phosphorus pentoxide will stand the highest temperatures
without decomposing. Nitric acid is a powerful oxidising
agent : phosphoric acid has no oxidising properties in any of
its various forms. The elements themselves, too, differ as
widely in their properties as any pair of elements. Nitrogen
will combine directly with oxygen only under very exceptional
circumstances, and will not combine directly with chlorine
at all. Phosphorus, on the other hand, combines with both of
these elements at the ordinary temperature. Owing to this
comparative inertness of nitrogen it is found free in nature ;
phosphorus, on account of its activity, is always found in
the combined state.

If, now, we compare phosphorus and sulphur together,

PHOSPHORUS 219

we find that although they belong to different groups of
elements and show no similarity in the formulae of their
compounds, they yet closely resemble each other in .many
points of their actual behaviour. Thus they are both solids,
comparatively easily fusible, and capable of existing in several
modifications. They are both combustible, and unite readily
with chlorine and the other halogens. Their hydrogen
compounds are both gaseous and give precipitates of phos-
phides or sulphides with many metallic salts which on
treatment with acids regenerate the original hydrogen com-
pounds. Sulphuric anhydride, SO3, as well as phosphoric
anhydride, P2O5, has a great tendency to combine with water,
a tendency not nearly so well developed in the lower oxides,
sulphurous anhydride, SO 2, and phosphorous anhydride,
P2O3. The chlorides, both of phosphorus and of sulphur,
are decomposed by contact with water.

The points of difference which should perhaps be chiefly
emphasised, are the feebly marked acid character of sul-
phuretted hydrogen as contrasted with the very feeble basic
properties of phosphuretted hydrogen, of which we have
evidence in the formation of phosphonium iodide, PH4I.
Sulphuric acid, again, can be easily reduced to sulphur
dioxide when it is not mixed with water; whilst phosphoric
acid cannot be reduced under any circumstances to form a
lower oxide.

CHAPTER XXIX

SILVER— COPPEB—MBBOUEY

The elements hitherto considered in the systematic portion of
this book form no basic oxides, and are known as the non-
metallic elements or non-metals. The metals, which have now
to be taken up, all form basic oxides, and can act as the
positive radicals of salts. Metals may or may not form acidic
oxides ; the common metals considered in the following pages
scarcely do so at all.

SILVER

Silver is a metal which occurs in considerable quantity
free in nature. It is generally found, however, not as the
metal, but in combination with sulphur as silver sulphide

The silver may be extracted from silver sulphide in many
ways, of which the following is an example. The crushed
silver ore is ground in mills with mercury and water, which
contains a little salt in solution. The silver sulphide is
slowly attacked by the mercury with formation of metallic
silver and mercuric sulphide, according to the equation —

AggS + Hg = HgS + 2Ag

Since a considerable excess of mercury is used, the silver
which is liberated dissolves in the liquid quicksilver and
forms what is called a silver amalgam. The process is
hence called an amalgamation process for extracting silver.
At the end of the reaction the mercury which contains the
silver is run off from the other products and subjected to
distillation in retorts. The mercury being volatile distils
off, and the metallic silver remains behind.

A large amount of silver is now extracted from argenti-

220

SILVER— COPPER— MERCURY 221

ferous lead. A great many lead ores, which consist chiefly
of lead sulphide, PbS, contain considerable quantities of
silver sulphide, and in the process of getting lead from lead
sulphide, the silver sulphide is at the same time converted
into metallic silver, which dissolves in the lead. Several
processes are in vogue for recovering this silver from the lead,
and an account of one of them will be given in the next
chapter.

Silver is a white metal, very soft and very tough, so that
it can be easily drawn into wire or beaten into foil. It is
unaffected by any of the atmospheric gases, except by the
sulphuretted hydrogen which is found in towns where coal
is burned. This gas tarnishes a silver surface owing to the
production of black silver sulphide.

Silver is so soft that it is practically impossible to employ
it in the pure state for the production of ornaments or coins.
In order to give it the requisite amount of hardness, it must
be alloyed or mixed with some other element, usually copper.
The ordinary silver coinage of Great Britain consists of

Silver 92.5%

Copper 7.5%

Pure metallic silver cannot be made to combine directly
with oxygen at any temperature, nor can it be made to
decompose water. Hydrochloric acid and dilute sulphuric
acid are also practically without action on it. It is easily
dissolved, however, by nitric acid and by concentrated
sulphuric acid with formation of silver nitrate and silver
sulphate respectively.

The chief soluble silver salt is silver nitrate, AgNOg,
which is also known under the name of lunar caustic. From
it most of the other silver compounds are prepared by
precipitation.

If we wish, for example, to prepare silver oxide, Ag20,
we can do so by adding a solution of sodium hydroxide
to a solution of silver nitrate. A brown precipitate falls
Qut, which is sometimes said to be silver hydroxide. It is
doubtful, however, if the brown substance is really the
hydroxide. At all events, it very readily loses water and

222 INORGANIC CHEMISTRY

becomes converted into silver oxide, the equations repre-
senting these reactions being —

AgNOa + NaOH = AgOH + NaNOg
2AgOH = AgjO + H2O

The silver oxide thus prepared parts with its oxygen on
heating, and becomes metallic silver according to the
equation —

2Ag20 = 4Ag + O2

When brought into contact with hydrogen peroxide (p. 169)
it decomposes in a similar manner, the equation in this case
being —

AggO + H2O2 = 2Ag + H2O + O2

The halogen salts of silver — namely, silver chloride^ silver
bromide^ and silver iodide — are all insoluble in water, and are
usually prepared by precipitating a solution of silver nitrate
by means of sodium chloride, bromide, or iodide. Silver
chloride is a pure white substance; the bromide and iodide
are pale yellow in appearance. All of these halogen salts
of silver are affected by hght, assuming a dark violet tint when
exposed to light for a sufficient length of time. Owing to their
sensibility to light they are employed in photography, the sensi-
tive substance in most photographic plates being silver bromide.

Although the soluble salts of silver are colourless, many of
the insoluble compounds of silver are coloured. Thus silver
oxide or hydroxide is brown, silver sulphide black, silver
iodide yellow, silver phosphate yellow, silver arsenate brown,
silver chromate crimson, and so on. On this account silver
salts, which are mostly insoluble, are very often used in
distinguishing the different salt radicals from each other.

COPPER

Metallic copper occurs free in nature in the neighbourhood of
Lake Superior, but most of the copper found in commerce is
produced either from copper sulphide CU2S, or from copper
pyrites CuFeSj, which is a double sulphide of copper and iron.
The ordinary process adopted in this country for the extrac-
tion of copper from copper pyrites is a somewhat complicated

SILVER—COPPER— MERCURY 223

one, and cannot be described here. There is, however, a
very simple wet process in use for extracting copper from
poor ores. When moist copper sulphide is left in the air,
the oxygen of the air converts the insoluble copper sulphide
into soluble copper sulphate. This may then be dissolved
out in water, and the solution may be made to yield metallic
copper by the addition of scrap iron according to the
equation —

CUSO4 + Fe = FeS04 + Cu

The precipitated copper is afterwards melted and refined.

Metallic copper in the pure or nearly pure state is -used
as a conductor of electricity. Copper, however, is mostly em-
ployed in the production of copper alloys. Thus brass con-
tains about two-thirds copper and one-third zinc; German
silver contains about two-thirds copper, the other components
being zinc and nickel ; bronze contains copper along with tin
and sometimes zinc ; our bronze coinage, for example, contains
95% of copper, with 4 of tin and i of zinc.

Copper is easily distinguished from other metals by its
warm red colour. Copper is not attacked at the ordinary
temperature by oxygen which is free from carbon dioxide, nor
is it attacked at any temperature by water ; but when exposed
for a long time to air, it becomes covered with a green incrusta-
tion of basic copper carbonate by the joint action of the oxygen,
moisture and carbon dioxide in the air. Although a com-
paratively soft metal, copper is considerably harder than silver.

It will not dissolve in hydrochloric or dilute sulphuric acid,
but when heated with concentrated sulphuric acid it pro-
duces copper sulphate and sulphur dioxide according to the
equation —

Cu -♦■ 2H2SO4 = CUSO4 + 2H2O -f SO2

At the same time some of the sulphuric acid is reduced still
further, so that the mixture rapidly becomes black from
formation of copper sulphide. Copper is readily attacked by
nitric acid, the copper being converted into copper nitrate,
and the nitric acid reduced for the most part to nitric
oxide, the equation being —

sCu -f SHNOs = 3Cu(N03)2 + 4H2O + 2NO
Copper forms two sets of salts, the cuprtc salts with the

224 INORGANIC CHEMISTRY

radical Cu", which are the ordinary salts of copper, and the
cuprous salts, with the radical Cu', which are comparatively
rarely met with.

The most common copper salt is bluestone or blue vitriol,
CuS04,sH20. This substance can be made by dissolving
copper in oil of vitriol, and crystallising the copper sulphate
thus produced from water; or it can also be made, as has
already been indicated above, by the action of atmospheric
oxygen on the natural sulphide. It crystallises in large blue
crystals, and is moderately soluble in water. When heated
to I GO*, it loses four-fifths of its water of crystallisation, but
only "parts with the remaining fifth when the temperature
rises to about 200'.

The nitrate can be prepared by dissolving copper in nitric
acid. The solution thus obtained deposits blue crystals with
the formula Cu(N03)2,3H2 0. These crystals are deliquescent,
and when heated lea^^ a black residue of cupric oxide CuO.

Gupric cUoride is easily prepared by heating cupric oxide
with hydrochloric acid, and allowing the solution to crystallise.
The chloride separates in the form of green crystals, having
the formula CuCl2,2H20.

A concentrated solution of cupric chloride is green, but when
diluted with water it becomes blue like the other soluble
cupnc salts.

Cuprous chloride, CuCl, can be readily procured by boiling
cupric chloride and concentrated hydrochloric acid with
metallic copper, the reaction which takes place being repre-
sented by the equation —

CUCI2 + Cu = 2CuCl

The solution obtained in this way is brown in colour, but
when poured into a large quantity of water it deposits a
white precipitate. This white precipitate consists of cuprous
chloride, which is insoluble in water, though soluble in con-
centrated hydrochloric acid. A solution of cuprous chloride
in hydrochloric acid is frequently employed in gas analysis
for absorbing carbon monoxide, which it readily dissolves.

Corresponding to the two sets of copper salts we have two
oxides of copper, cupric oxide, CuO, and cuprous oxide, CugO.
We have already seen that cupric oxide can be made by heating
copper nitrate. Cuprous oxide is easily produced as a bright

SILVER— COPPER— MERCURY 225

red precipitate when a solution of copper sulphate is boiled
with glucose and excess of caustic alkali. The glucose serves
to reduce the copper from the cupric to the cuprous state.
When copper is heated in the air both of these oxides are
formed. The ordinary copper scale produced on copper in
this way is cupric oxide on the outside and. cuprous oxide in
the interior.

When sodium hydroxide is added to a soluble cupric salt, a
blue precipitate of cupric hydroxide is obtained according to
the equation —

CUSO4 + 2NaOH - Cu(0H)2 + Na2S04

When the liquid containing this precipitate is heated to the
boiling point, the cupric hydroxide loses water, and becomes
black from formation of cupric oxide —

Cu(0H)2 = CuO + HgO

Basic carboiiates of copper are found as minerals in nature.
They have either a blue or green colour, the most valuable
varieties being known as malachite and azurite. The blue
precipitate obtained by adding sodium carbonate to copper
sulphate solution is also a basic carbonate.

Cuprous sulpMde, CU2S, occurs in large quantity in nature,
and is one of the chief ores of copper. Cupric sulphide, CuS,
which also occurs in nature, but in much smaller quantity, can
be easily prepared in the laboratory as a black precipitate by
passing sulphuretted hydrogen through a solution of a cupric
salt —

CUSO4 + HgS = CuS + H2SO4

MERCURY

Mercury or quicksilver, which presents absolutely no physi-
cal resemblance to metallic copper, is yet very similar to copper
in many of its chemical actions. Metallic copper is a red,
moderately hard metal of high melting point; mercury is
a pure white metal, which is liquid even at the ordinary tem-
perature. Mercury, like copper, occurs in small quantity in the
metallic state, but it is chiefly found in the form of mercuric
sulphide, HgS, which is commonly known as cinnabar. The
p

d26 INORGANIC CHEMISTRY

extraction of metallic mercury from cinnabar is a very simple
affair. The mercuric sulphide is roasted in kilns in an adequate

supply of air. The sulphur
is converted into gaseous sul-

Fig. 38.-Aludels for Condensing P^^"' d'°'^<*^ ^^^^^ ^he metallic
Mercury. mercury comes off as vapour

These aludels are earthenware vessels, ^t the high temperature gCttC-
which are used in long series, the stem of rated by the COmbuStion. The
each fitting into the neck of its successor. ^ • i j j

mercury vapour is cooled and
condensed in large chambers or in aludels (fig. 38) to
liquid mercury, which may then be purified, first by squeezing
through wash leather, and, finally, by distillation.

Mercury, which, when sufficiently cooled, is a hard crystalline
metal, melts at -39", and boils at 360°, the molecular formula
of the vapour being Hg. It is used, on account of its
being a liquid through so great a range of temperature,
for filling thermometers, and also in the construction of
barometers and many other scientific instruments. Its
chief practical use, however, is in the extraction of silver and
gold by amalgamation processes, and in the silvering of
mirrors. The " silver " on the back of mirrors is really an alloy
of tin and mercury, or tin amalgam, the alloys of mercury
being always called amalgams.

Mercury is not attacked by air at the ordinary temperature,
and retains its bright surface permanently. It has already
been stated that when heated in air for a long time at its
boiling point it gradually combines with oxygen to produce
mercuric oxide, HgO, which, however, easily splits up again
into its elements on further heating. Mercury does not de-
compose water at any temperature. It does not dissolve
in hydrochloric acid, or in dilute sulphuric acid. Like
copper, it forms two sets of salts, mercurous salts, with the
radical Hg', and mercuric salts, with the radical Hg*\

When heated with concentrated sulphuric acid, it behaves
like copper, forming mercuric sulphate and sulphur dioxide,
according to the equation —

Hg + 2H2SO4 = HgS04 + 2H2O + SOa
It also dissolves readily in nitric acid. When the nitric

SILVER— COPPER— MERCURY 227

acid is warm and dilute, mercuroos nitrate is produced in
accordance with the equation —

3Hg + 4HNO3 = 3HgN03 + NO + 2H2O

When the nitric acid is more concentrated, mercuric nitrate is
produced according to the equation —

3Hg + 8HNO3 = 3Hg(N03)2 + 2NO + 4H2O

The dilute nitric acid thus oxidises the metallic mercury
to the mercurous state of oxidation, whilst the hot concentrated
nitric acid oxidises it further to the mercuric stage.

Both the nitrates of mercury are easily soluble in water
containing a little nitric acid, but when they are treated with
pure water, they decompose with formation of a quantity of
insoluble basic nitrates and liberation of nitric acid. The
solutions of mercuric or mercurous nitrate, therefore, which
are used in the laboratory always contain a considerable
quantity of free nitric acid.

The two chlorides of mercury are the commonest mercury
compounds. Mercuric chloride, HgCl2, is prepared by dissolv-
ing metallic mercury in concentrated sulphuric acid, evaporating
the mercuric sulphate thus obtained to dryness, and then
distilling it with common salt. Double decomposition takes
place according to the equation—

HgS04 + 2NaCl = HgClg + NagSO^

and the mercuric chloride, being the most volatile of the
substances concerned in this reaction, distils off, leaving a
residue of sodium sulphate. The mercuric chloride vapour
condenses to a crystalline mass which can readily be purified
further by sublimation. On account of the ease with which
mercuric chloride can be sublimed it is frequently known by
the name of corrosive sublimate.

Mercuric chloride is easily soluble in boiling water, but not
very soluble in cold water. It is a very poisonous substance,
and can only be administered as a drug in small doses.

Mercurous chloride, HgOl (or, as it is sometimes written,
Hg2Cl2), is prepared by adding metallic mercury to the

228 INORGANIC CHEMISTRY

mixture of common salt and mercuric sulphate before sub-
liming.

HgS04 + 2NaCl + Hg = 2HgCl + Na2S04

When mercuric chloride and mercury are vaporised together,
the vapours on cooling combine to produce mercurous
chloride, which again is largely dissociated into mercuric
chloride and mercury on resublimation.

HgCla + Hg <t 2HgCl

Mercurous chloride or calomel produced in this way is usually in
the form of an amorphous powder which is practically insoluble
in water and in dilute acids. Like mercuric chloride, calomel
is also used in medicine, but on account of its not being nearly
so soluble it may be administered in much larger doses.

When sodium hydroxide is added to a soluble mercuric salt,
a yellowish precipitate of mercuric hydroxide is obtained, which
on drying leaves a red residue of mercuric oxide, HgO. This
oxide, when heated, decomposes, as we have already seen, into
metallic mercury and oxygen. When sodium hydroxide is
added to a solution of a mercurous salt, a black precipitate is
obtained which is supposed to consist of mercurous oxide,
Hg20. This substance, however, if it exists, is very readily
decomposed into mercuric oxide and mercury.

Hg^O = HgO + Hg

Mercuric sulphide, HgS, is not only found in nature as
the red mineral cinnabar^ but is manufactured on a large scale
and used as a red pigment under the name of vermilion.
When precipitated from a solution of a mercuric salt it is
black.

Mercury salts behave in a peculiar way when treated with
a solution of ammonium hydroxide. In general, if a hydroxide
is precipitated from a solution of a metallic salt by sodium
hydroxide, the same precipitate is obtained when ammonium
hydroxide is used as precipitant. In the case of mercury,
however, it is quite different. If a solution of ammonia is
added to a mercuric salt, a white precipitate is produced,
instead of the reddish yellow precipitate which is obtained with
sodium hydroxide. This white precipitate contains mercury.

SILVER— COPPER— MERCURY 229

but no longer as the positive radical. The mercuric mercury
has replaced some of the hydrogen in the ammonium radical,
so that the white precipitate contains a complex positive radical
into which mercury, nitrogen, and hydrogen enter. The
formula of the chloride is usually represented as (NH2Hg)*Cr.
When ammonia solution is added to a mercurous salt, a black
precipitate is formed. This precipitate is not mercurous oxide
or hydroxide, but again contains a complex ammoniacal positive
radical.

The oxidation of a mercurous to a mercuric salt can easily
be effected by strong nitric acid, and a mercuric salt can be
easily reduced to a mercurous salt by means of stannous
chloride.

2HgCl2 + SnClg = 2HgCl + SnCU

This reaction is made use of in testing for mercury. When
stannous chloride is added to the solution of the mercuric
salt, a white precipitate of mercurous chloride is immediately
obtained, which may afterwards become grey by further
reduction to minute globules of metallic mercury.

It will be seen from what has preceded, that there is con-
siderable resemblance between silver, copper, and mercury in
their chemical properties. The metals themselves have no
action on water at any temperature, and are not attacked by
hydrochloric or dilute sulphuric acid. They all dissolve, how-
ever, in concentrated sulphuric acid and in nitric acid. Silver
forms only one set of salts, copper and mercury each forms two
sets ; and the similarity that there is between silver compounds
on the one hand, and copper or mercury compounds on the
other, exists between the silver salts and the less oxidised salts
of the copper or mercury. Thus, while cupric chloride and
mercuric chloride are soluble in water, cuprous chloride and
mercurous chloride are, like silver chloride, quite insoluble in
water. The same holds good for many other mercurous and
cuprous compounds.

CHAPTER XXX
LEAD— TIN

LEAD

Metallic lead occurs in excessively small quantities in nature.
It is chiefly met with when in the form of sulphide, and
occasionally in the form of carbonate or sulphate.

The mineral galena or lead sulphide, PbS, is by far the
most important ore of lead, and nearly all commercial lead is
prepared from it. The preparation of metallic lead from lead

fimmmmmm

Fig. 39. — Reverberatory Furnace for Lead-smelting.

In a reverberatory furnace the fuel is burned on a grate G, and is
separated from the bed of the furnace J9, on which the chemical pro-
cess takes place, by a bridge R. The flames and hot gases from the
fire strike the roof A and are reflected or " reverberated" back on the
bed of the furnace, which in this case is hollowed out to collect the
molten lead. The gases pass off by the flue F. // is a. hopper to
introduce the ore, here lead sulphide.

sulphide consists first in the partial oxidation in a reverberatory
furnace of the lead sulphide to lead oxide and lead sulphate.
These compounds are then reduced by fusing them with the
unoxidised lead sulphide, which contains the reducing element
sulphur. The equations for the actions are as follows. First
there are the two oxidising equations —

PbS + 2O2 = PbS04
2PbS + 3O2 = 2PbO + 2SO0

230

LEAD— TIN 231

Then there are the two reducing equations —

PbS04 + PbS = 2Pb + 2S0g
2PbO + PbS = 3Pb + SO2

Sulphur dioxide is the only other product besides metallic
lead.

It was stated under the heading of silver that most lead ores
contain a small quantity of silver ores, and that the silver
ultimately finds its way into the metallic lead. It is profitable
in many cases to extract the silver from lead, and various pro-
cesses have been devised for this extraction.

One of the simplest of these is Pattinsan^s process^ which
takes advantage of the fact that the first solid which separates
out from a molten mixture of lead and silver is pure lead. If,
therefore, a considerable proportion of the molten argentiferous
lead is allowed to solidify, and is then removed, nearly all the
silver remains in the liquid residue. By systematically carrying
out this separation the silver may be concentrated in a com-
paratively small quantity of lead. This rich lead is then
subjected to a process called cupellation. Cupellation consists
in the oxidation of metallic lead to lead oxide by heating
the molten lead to bright redness in air. The skin of oxide
which forms on the surface of the molten metal is soaked up
by a porous bone-ash hearth or cupel^ in which the molten
lead is contained. A fresh surface of lead is thus exposed, and
this in turn becomes oxidised. The oxide is again absorbed :
and so the process goes on until all the lead has been removed,
and the metal which remains is pure silver.

Metallic lead is chiefly useful on account of its softness
and easy fusibility. It can be readily formed into wire or pipe
by squirting the soft metal through a steel die, and can also be
rolled into sheet lead, which may be beaten into the shape of
any object which it may be desirable to cover or line with it.
Metallic lead does not decompose water alone, although it
does so in presence of oxygen ; and is not attacked by hydro-
chloric acid or dilute sulphuric acid. Even concentrated
sulphuric acid attacks metallic lead only slowly, so that lead
vessels rfe very often employed on the large scale in operations
where sulphuric acid is used. Nitric acid dissolves lead very
readily with production pf lead nitrat^i

232 INORGANIC CHEMISTRY

Lead nitrate, Pb(N0s)2, and lead acetate, Fb(C2H302)2,
are the commonest soluble salts of lead. The former may
be made by dissolving metallic lead in nitric acid, and the
latter is made by dissolving lead monoxide in acetic acid.

When sodium hydroxide is added to a soluble lead salt a
white precipitate of lead hydroxide is produced.

Pb(N03)2 + 2NaOH = Pb(0H)2 + 2NaN03

This precipitate when heated loses water, and a residue of
lead monoxide or litharge, PbO, remains. This oxide can
be produced directly by heating lead in air. It has no
tendency to give up oxygen on further heating, but rather
absorbs oxygen from the air and is converted into the oxide
PbaO^.

6PbO + 02 = aPbgOA

This oxide is an important one, for when properly prepared
it forms the pigment known as red lead.

Red lead, Pb304, is not a basic oxide, and no series of salts
corresponds to it. When it is treated with nitric acid it
behaves as if it were a mixture of lead monoxide, PbO, and
lead dioxide, PbO 2. The lead monoxide dissolves in the
nitric acid with production of lead nitrate —

PbO + 2HNO3 = Pb(N0s)2 + H2O

and the lead dioxide, or lead peroxide as it is usually called,
remains as a black residue.

Lead dioxide, PbO 2, has very feebly basic properties.
Owing to its reluctance to act as a basic oxide, lead dioxide
does not dissolve readily in acids, and in particular is quite
insoluble in nitric acid. When warmed with hydrochloric
acid it acts like manganese dioxide (p. 170), yielding lead
chloride and chlorine.

When hydrochloric acid or a soluble chloride is added to
a soluble salt of lead, a white precipitate of lead chloride,
PbCl2, is immediately formed. This chloride is very sparingly
soluble in cold water, but dissolves readily in boiling water,
from which it is re - precipitated in crystalline scales on
-cooling.

When a soluble sulphate is added to a lead salt, a white
precipitate of lead sulphate, PbS04, is obtained. It is almost

LEAD— TIN 233

insoluble in water, but dissolves in concentrated sulphuric
acid.

Normal lead carbonate, PbCOa, occurs in nature as the
mineral cerussite, and can also be prepared by precipitating a
solution of lead acetate with ammonium carbonate. If sodium
carbonate is used to precipitate a solution of a lead salt, a
basic carbonate is produced, the composition of whieh varies.
A basic carbonate of lead having the composition 2PbC03,
Pb(0H)2, is very extensively used as a pigment, and is known
as white lead.

Lead sulphide, PbS, not only occurs as galena, but can
easily be formed by precipitating any soluble lead salt with
sulphuretted hydrogen. Both the natural and precipitated
varieties are black.

TIN

Metallic tin is scarcely ever met with in the free state in
nature. It almost invariably occurs oxidised in the form of
tin-stone^ which is more or less pure tin dioxide, Sn02.

After a mechanical treatment and a preliminary roasting in
air to get rid of impurities, the tin-stone is heated in a
reverberatory furnace along with powdered coal, which takes
the oxygen of the tin dioxide and liberates metallic tin, the
equations for the action being —

SnOa

-h

—

2CO

SnOg

CO2

Tin is a pure white metal which generally has a marked
crystalline structure. At the ordinary temperature it is very
malleable, and can easily be rolled into the thin sheet known
as tin-foil. Its melting point is only a little above 200**.
Metallic tin resists the action of air at the ordinary tempera-
ture, but it can scarcely be used in the pure state for the
manufacture of utensils or boxes on account of its compara-
tively high cost. When a perfectly clean plate of iron, how-
ever, is dipped into molten tin, the tin will adhere to it, and
cover it with a uniform coating, which protects it from the
action of the air which otherwise would rust the iron. This
iron, with a covering of tin, is what we know as tin-plate^
and is the material of which " tin " cans, etc., are made.

234 INORGANIC CHEMISTRY

Tin is a constituent of many alloys — such as solder and
pewter^ which are alloys of tin and lead, — and gun metal
or bronze^ which are alloys of tin and copper.

Tin, although not acted upon by air at the ordinary
temperature, combines readily enough with oxygen when
heated. Thus, if a piece of tin-foil is held in a bunsen
flame it will melt, and at the same time oxidise with pro-
duction of copious white fumes of tin dioxide, the equation
for the combustion of tin being —

Sn -h O2 = SnOa

Tin does not decompose water at a boiling heat, but if
steam is passed over red-hot metallic tin, hydrogen is pro-
duced according to the equation —

Sn + 2H2O = Sn02 + 2H2

In accordance with this power to decompose water, metallic
tin decomposes hydrochloric and sulphuric acids at the
ordinary temperature with production of hydrogen.

Sn + 2HCI = SnCl2 + Hg

Tin is not dissolved by nitric acid, but is oxidised to an
insoluble white powder called metastannic acid, the equation
for the formation of which is —

3Sn + 4HNO3 -h H2O = 3H2Sn03 + 4NO

The chief soluble compounds of tin are stannous chloride
SnCl2, and stannic chloride SnCl4.

Stannous chloride^ SnCl2^ is prepared by dissolving tin in
hydrochloric acid, the chloride crystallising out when the con-
centrated solution is cooled. Stannous chloride is a reducing
agent, and a solution of it on exposure to the air absorbs
oxygen. This oxidation may be prevented by adding some
hydrochloric acid to the solution and keeping it in contact
with metallic tin.

Examples of the reducing action of stannous chloride have
already been given (p. 103), the reduction of mercuric chloride
to mercurous chloride (p. 229) being the most characteristic,
and one frequently used as a test for tin in the form of
stannous salt

LEAD--TIN 235

Pure stannic chloride^ SnCl4^ may be obtained by heating
metallic tin in a current of dry chlorine, the formation taking
place according to the equation —

Sn + 2CI2 = SnCU

A solution of stannic chloride is prepared by heating metallic
tin with a mixture of hydrochloric and nitric acids, the nitric
acid serving to oxidise the tin from the stannous to the
stannic stage of oxidation.

Stannic chloride has none of the characteristics of a salt.
It is a liquid boiling at a temperature not much above that
of the boiling point of water, and is easily decomposed by
alkalies and even water, with formation of the corresponding
hydroxide.

SnCU + 4H2O = Sn{0H)4 + 4HCI

This hydroxide is not a base but rather a very feeble acid,
being generally known as stannic acid.

Stannous sulphide^ SnS, and stannic sulphide, SnSg^ may
be formed by the action of sulphuretted hydrogen on the
corresponding chlorides, thus —

SnClg + H2S = SnS + 2HCI

SnCU + 2H2S = SnS2 + 4HCI

The former is a brown and the latter a yellow precipitate.
Stannic sulphide is prepared in the dry way in the form of
golden yellow scales called " mosaic gold."

CHAPTER XXXI
ZINC— ALUMINIUM

ZINC

Zinc does not occur in nature in the metallic state, but occurs
chiefly as zinc blende, which is the sulphide ZnS, or as calamine,
which is the carbonate ZnCOg. The production of metallic
zinc from these ores is very simple. The sulphide and
carbonate are first heated strongly in the air, and converted
into zinc oxide, the equations for the roasting being —

2ZnS + . 3O2 = 2ZnO + 2SO2

ZnCOs = ZnO + CO2

The zinc oxide obtained in this way is then heated in long fire-
clay retorts, together with car-
bon in the form of coal. The
carbon reduces the zinc oxide

Fig. 40.— Retort for Reduction to metallic zinc —
of Zinc Oxide.

The oxide and carbon are heated to- ZnO -f- C = Zn -H CO

gether in the fireclay retort A^ which is
attached by means of the adapter B to the

iron condenser C, which projects beyond ^hich, at the temperature of
the furnace used for heating the retorts. ' . . i • i -i

the reaction, is above its boil-
ing point. The zinc, therefore, comes off as vapour, and is con-
densed in iron tubes attached to the open end of the long
fireclay retorts.

Zinc is a metal which in many ways resembles tin. It can
be distinguished from tin, however, by its bluish colour. It
melts below a red heat, and at the ordinary temperature is
not attacked by air or by water. When heated in air it
burns with production of zinc oxide ZnO. This may be
seen by holding a strip of zinc foil in the bunsen flame,
the zinc oxide which is produced appearing as copious
white fumes.

236

ZINC— ALUMINIUM 237

When water in the form of steam is passed over heated
zinc, it is decomposed with evolution of hydrogen.

Zn + H^O = ZnO + H

Zinc, like tin, is often used to form a protective coating for
iron. The iron to be protected is carefully cleaned and then
dipped into a bath of molten zinc. The zinc adheres to the
clean iron surface and protects it from the action of the air.
Iron which has been treated in this way is called galvanised
iron, although no galvanic action is used in its production.

Zinc dissolves readily in all acids, hydrogen being evolved
with hydrochloric acid or dilute sulphuric acid. The
soluble salts of zinc can thus be easily made by dissolving
the metal in the appropriate acid. The • most common of
soluble zinc salts is zinc sulphate or zinc vitriol, ZnS04,7H20.
This substance may be made on the large scale by roasting
zinc sulphide at a carefully regulated temperature. The
sulphide combines with oxygen, according to the equation —

ZnS + 2O2 = ZnS04

and the zinc sulphate thus produced may be dissolved in
water and purified by recrystallisation.

When sodium hydroxide is added to a solution of zinc
sulphate, a white precipitate of zinc hydroxide is produced.

ZnSO^ + 2NaOH = Zn(0H)2 + Na2S04

This hydroxide readily loses water on heating, and is con-
verted into zinc oxide.

Zn(0H)2 = ZnO + HgO

Zinc oxide produced in this way, or by burning zinc in
air, which is the method adopted for its commercial produc-
tion, is a pure white substance which has the characteristic
property of becoming bright yellow on heating. On being
cooled it regains its pure white colour. It is employed as
a pigment under the name of zinc white.

When a solution of sodium carbonate is added to a soluble
zinc salt, a white precipitate of basic zinc carbonate is
produced.

238 INORGANIC CHEMISTRY

Although the zinc blende which occurs in nature is almost
invariably coloured black by admixture with ferrous sulphide,
pure zinc sulphide is colourless. It may be produced as a
white precipitate by adding a solution of an alkaline sulphide
to a soluble zinc salt.

ZnS04 + NaaS = ZnS + NagSO^

Zinc chloride, ZnClg, is prepared by dissolving zinc oxide or
metallic zinc in hydrochloric acid. Its concentrated solution
has a very caustic action, and will dissolve paper or cotton.
It is usually cast in the form of sticks, which have a remark-
able attraction for water,- and are frequently used for removing
traces of water in certain chemical actions.

ALUMINIUM

Aluminium is a metal which never occurs in the free state
in nature. It is always found in the oxidised condition,
chiefly in the form of silicate. China clay or kaolin is a
very nearly pure hydrogen aluminium silicate. Ordinary clay
consists chiefly of aluminium silicate, but contains besides
silicates of iron and other metals which give it its colour.
Fire clay consists of ordinary clay incorporated with silica.
Clays which contain much iron have a red colour ; china clay
is colourless ; while pipe clay contains very little iron and is
almost without colour.

AiiiTniTiiiim oxide or alumina, AI2O3, is a very hard sub-
stance, the pure mineral being termed corundum. An
impure variety of corundum is largely used on account of its
hardness as a polishing powder under the name of emery.
Some precious stones, such as sapphire and ruby^ consist of
nearly pure alumina. A mineral known as bauxite^ which
contains about two-thirds of its weight of alumina, has recently
attained importance as a source of the metal aluminium.

In order to prepare metallic aluminium from bauxite, the
mineral is first fused with sodium carbonate, when sodium
aluminate is formed, according to the equation —

sNaaCOa + Al^Og = 2Na8A103 + 3CO2

Sodium aluminate

This being a sodium salt is soluble in water, and the

ZINC— ALUMINIUM 239

aluminium may thus be removed from iron and other im-
purities which remain behind in the form of insoluble oxides.
By passing carbon dioxide into the solution of sodium
aluminate, the aluminium may be precipitated as aluminium
hydroxide, decomposition occurring according to the equation —

aNaaAlOg -I- 3CO2 + 3H2O = 2A1(0H)3 + 3Na2C08

This aluminium hydroxide is then heated to convert it into
oxide —

2A1(0H)3 = AlaOa + 3H2O

When the aluminium oxide thus obtained is fused in an
electric furnace, and a current of electricity passed through
the fused material, electrolysis takes place, metallic aluminium
being liberated at the kathode, and oxygen being liberated at
the anode. At the high temperature of the electric furnace
the metallic aluminium is liquid and can be drawn off from
time to time.

Aluminium is a pure white metal which resists the action of
air and water at the ordinary temperature. This property, along
with its strength and low specific gravity (which is only
about one-third of that of iron), renders it very useful, and it
is now being used for the manufacture of many metallic
articles in which lightness combined with strength is desired.

It does not decompose water at the ordinary temperature,
except in conjunction with mercury (p. 166), but will decom-
pose steam when highly heated. It is scarcely attacked by
nitric acid, and is not easily attacked by dilute sulphuric acid,
but it dissolves readily in hydrochloric acid, according to the
equation —

2AI + 6HCI = 2AICI3 + 3H2

When aluminium is heated to a high temperature in the air it
burns, and is converted into aluminium oxide AlgOg.

The commonest compound of aluminium is alum, which is
a double sulphate of aluminium and potassium or ammonium,
the formulae of the compounds being respectively —

K2S04,Al2(S04)3,24H20 (NH02SO4,Al2(SO4)3,24H2O
Potassium alum Ammonium alum

Ordinary alum may be either of these compounds or a
mixture of both.

240 INORGANIC CHEMISTRY

Alum may be prepared from a silicate of aluminium, say
pipe clay, in the following manner : — The clay is ground to a
fine powder and heated with concentrated sulphuric acid until
it forms a paste, which is then exposed to the air for some
weeks. The sulphuric acid attacks the aluminium silicate,

with formation of silicic acid HgSiOg,
and aluminium sulphate Al2(S04)3.
This aluminium sulphate is soluble,
and may be separated from the other
materials, which are insoluble, by treat-
ing the mass with water. When the
solution is evaporated, aluminium sul-
phate crystallises out, but is not easy

Fig. 4i.-Crystal of ^^ P""^^ by crystallisation. It is, there-
Alum, fore, mostly dissolved up again, and to
The alums crystallise in the solution is added either potassium

'^'J°Znf.^iiH^^'^ ?>• ammonium sulphate, when the spar-
solid, with all ihe faces ingly solublc alums Separate out and can

equilateral triangles. j*i i_ 'i^ j i_ ^ ii- ^*

readily be punned by recrystallisation.
Both alum and aluminium sulphate are extensively used in
dyeing and paper-making.

When a soluble hydroxide is added to a solution of an
aluminium salt, aluminium hydroxide is precipitated, the equa-
tion being —

Al2(S04)3 + 6NaOH = sNa^SO^ + 2A1(0H)3

This aluminium hydroxide can behave in two ways. With
strong acids it behaves as a base. Thus, it will dissolve in
hydrochloric acid with formation of aluminium chloride, as
follows : —

Al(0H)3 + 3HCI = AICI3 + 3H2O

But with strong bases it can also behave as an acid. Thus,
if excess of caustic soda is used in precipitating the solution
of an aluminium salt, the aluminium hydroxide first formed
dissolves up in the excess of sodium hydroxide, with produc-
tion of the soluble sodium aluminate, the equation for this
action being —

A1(0H)3 + 3NaOH = NagAlOs + sH^O
Here the aluminium hydroxide acts as an acid with respect to

ZINC— ALUMINIUM 241

the sodium hydroxide. It is only with regard to strong bases,
however, that the aluminium will act in this way. If, for
example, we take ammonia instead of caustic soda, we find
that the precipitated aluminium hydroxide is much less easily
dissolved up again. That is on account of ammonium hy-
droxide being so much weaker a base than sodium hydroxide,
that it is unable to remain permanently combined with the
very feeble acid aluminium hydroxide, unless it is present in a
very large excess so as to make up by its quantity for its lack
of strength.

Aluminium hydroxide is not only very feeble as an acid, it
is also very feeble as a base. Although its salts with strong
acids are stable enough, being only slightly hydrolysed in
solution (p. 141), aluminium acetate when boiled with
water is decomposed, the whole of the aluminium being
precipitated as a basic acetate and the weak acetic acid being
liberated. With an acid so weak as carbonic acid, aluminium
hydroxide can form no salt at all. Thus, when sodium
carbonate is added to a solution of aluminium sulphate,
although a white precipitate is produced, this precipitate is
not aluminium carbonate, as we might expect : it is aluminium
hydroxide. We may, if we choose, imagine that aluminium
carbonate is first produced from the aluminium sulphate and
sodium carbonate, according to the equation —

Al2(S04)3 + sNaaCOa = 3Na2S04 + Al2(C03)3

but that this aluminium carbonate, being a compound of a
very weak base with a very weak acid, is at once decomposed
by water with formation of aluminium hydroxide and carbonic
acid —

Al2(C03)3 + 6H2O = 2A1(0H)3 + 3H2CO3

Sulphuretted hydrogen like carbonic acid is a very weak acid.
Although it is possible to prepare aluminium sulphide in
the dry way, the product is at once decomposed by wafer with
formation of aluminium hydroxide and liberation of hydrogen
sulphide —

AI2S3 + 6H2O = 2A1(0H)8 + 3H2S

The precipitate which separates when sodium or ammonium
sulphide is added to a soluble aluminium salt is thus not
aluminium sulphide, but aluminium hydroxide,

CHAPTER XXXII

IBON

Iron occurs in the metallic state only in very small quantity,
chiefly in meteorites, which are not of terrestrial origin at all.
It is found combined very abundantly with oxygen and with
sulphur. The commonest compound with sulphur is iron
pyrites FeSg, which, although a convenient source of sulphur,
is not to be classed amongst iron ores — i.e. minerals from
which metallic iron can be profitably extracted. The chief
ores of iron are oxides and the carbonate, which occur in
a more or less pure condition, their composition being indi-
cated in the following table : —

Ferrous carbonate (spathic iron ore) . . . FeCOg

Ferric oxide (red haematite) . . . . FcgOg

Ferric hydroxide (in brown haematite) . . Fe(OH)3

Ferroso-ferric oxide (magnetic iron ore, lode-stone) Fe304

Metallic iron can be obtained from these ores by first of all
roasting them to get rid of carbon dioxide and certain im-
purities, and then reducing them in a blast furnace by means
of carbon in the form of coke, or of charcoal, if a very pure
iron is desired. The actions which go on in a blast furnace
are very complicated, but it would appear that the carbon
dioxide which is produced by the union of the oxygen of. the
hot air blast which is blown in at the bottom of the furnace
with the carbon of the fuel, is reduced, in the higher portions
of the furnace, to carbon monoxide, which in its turn reduces
the ferric oxide to metallic iron. This reduction occurs in a
comparatively cold part of the furnace, the temperature being
insufficient to melt the pure metallic iron, w^hich is one of the
least fusible of metals. The metallic iron, however, is capable
of taking up carbon, partly to form a carbide of iron, and the
resulting product, known as cast iron^ is easily fusible and sinks

242

IRON 243

to the bottom of the furnace, from which it can be withdrawn
as required. At the same
time the silica and Ume
found associated with the
iron compounds in the ores,
or specially added along
with them, unite Whether
to form a fusible slag, con- *
taining chiefly calcium sili-
cate, which lies as a layer on
the surface of the fused cast
iron and protects it from
oxidation by the blast.

Oast iron as it leaves the
furnace is made to run into
moulds and is there allowed
to solidify in the form of
bars. It consists of iron

with from 2 to 5 per cent. „

of carbon, partly as car-
bide, and partly crystallised ,
throughout the bar in the

form of graphite. When - Fig. 42.— Blast Furnace.

rflst imn i*i iTPatpH with ^" ^ blast furnace ihe fufl and ore are

■A iu "^^^'^ '*"" n,Lxed tog«l«, in Ih. boly o! .h= fpr-a«.

acid, the iron dissolves and imd not sepatawd n in a reverbcratory fur-
hydrogen is given off. This "h"u|h a'^Hef of'"b« /-"'ihi'tolt™ !rf
hydrogen, however, is by no 'ii= fumace 10 luppiy ih* wygto nectsswy

means pure, as it contains isclMelbyllieeoneC.whilSBloIiwdwSen
quantities of hydrocarbons ? '"^ cbarge Was to be introduce. The
2 J e .1. furnace gases escape by the flue F. Tb*

derived from the iron car- moiien cast-iron /caiiecis on tbc heutb of
bide, j«st .s the hydroc- SJ-SfAS-S ™ilfeS.".i'

bon acetylene is derived ore, which protects U from oiidation by tbe

from calcium carbide. There "" '"''

remains also a black residue which does not dissolve in acids,
and consists chiefly of carbon in the form of graphite. Cast
iron, as its name implies, can be cast in a mould to any
desired form. This is on account of the comparatively low
temperature at which it fuses, and also on account of its
expanding slightly on solidification, so that it enters the
smallest crevices of the mould and reproduces the details
exactly.

244 INORGANIC CHEMISTRY

Wrought iron or malleable iron may be produced from
cast iron by removing its carbon. This may be effected in
various ways, but in principle they are all the same. If the
cast iron is melted and exposed to air, the oxygen of the air
combines with the carbon in preference to the iron, so that
at a high temperature the whole or nearly the whole of the
carbon may be removed in the form of gaseous oxides of
carbon, a pasty mass of wrought iron remaining behind.
Wrought iron is much less easily fused than cast iron and
cannot be used for making casts. It is worked into shape
by being rolled or hammered to the desired form when hot.
It is not brittle like cast iron, but tough and fibrous.

Steel is intermediate in composition between cast iron and
wrought iron. It contains about ^ per cent of carbon, but
the quantity may vary according to the quality of steel which
is desired. Mild steel contains a comparatively small quantity
of carbon, and approximates in character to wrought iron.
Tool steel contains a larger proportion of carbon, and is more
like cast iron in its properties. The most valuable property
of steel is its capability of being tempered. If a piece of steel
is heated to redness and is then suddenly chilled in water or
oil, it becomes extremely hard, but is at the same time very
brittle. The hardness and brittleness may be removed to any
required degree by heating the steel to moderate temperatures
and allowing it to cool. If the steel is heated only slightly it
loses very little of its hardness and brittleness. If it is heated
to a higher temperature it becomes tough and elastic, and at
intermediate temperatures it may be made to assume inter-
mediate properties. The hardest steel easily scratches glass
and is excessively brittle, but the same steel may be tempered
down until it acquires the elasticity and toughness of a watch
spring. The temper of a knife-blade is intermediate between
these two extremes.

Steel of any desired composition may be made by melting
up cast iron and malleable iron in the proper proportions, and
most of the ordinary processes for the production of steel are
varieties of this method of manufacture.

Pure iron may be most easily prepared by electrolysis, or by
reducing an oxide of iron by means of aluminium, the equation
in the latter case being —

FejjOa + 2AI = AlgOs + 2Fe

IRON 245

Pure iron has a very high melting point, the temperature
being estimated at about i6oo'*«

Iron is not attacked by oxygen or air which is free from
moisture or carbon dioxide. In ordinary air, however, which
contains both of these substances, iron is slowly attacked at
the atmospheric temperature with formation of rust, which is
chiefly ferric hydroxide Fe(0H)3. The action of rusting
seems to consist in water and carbon dioxide jointly attacking
the iron, with production of ferrous carbonate FeCOs, which
is afterwards oxidised to ferric hydroxide Fe(0H)3 with
liberation of carbon dioxide. Iron does not attack water at
the ordinary temperature, but if steam is passed over red-hot
iron, the iron is oxidised, and hydrogen is set free. The
equation usually given for this reaction is —

3Fe + 4H2O = FegO^ + 4H2

Ferroso-ferric oxide

but it is doubtfub if the oxide actually formed has the com-
position here represented.

Iron is readily attacked by dilute sulphuric, hydrochloric,
and nitric acids, but it presents a peculiarity in resisting the
action of concentrated sulphuric and concentrated nitric acid.
No satisfactory explanation of this behaviour has yet been
arrived at.

Iron forms two sets of salts, the ferrous salts containing the
radical Fe", and the ferric salts containing the radical Fe*".
The commonest soluble ferrous salt is ferrous sulphate in the
form of green vitriol or copperas FeS04,7H20. This is
usually manufactured on the large scale by the slow atmos-
pheric oxidation of moist iron pyrites FeS 2 • The soluble ferrous
sulphate is extracted by water and recrystallised. It can
also be readily obtained by dissolving iron in diluted sulphuric
acid, and crystallising the solution. Ferrous sulphate, like all
other soluble ferrous salts, tends to take up oxygen from the
air, especially when in solution, and become oxidised to a
ferric salt. A solution of ferrous sulphate, therefore, always
contains some ferric sulphate, unless when freshly prepared.
Ferrous sulphate, on account of the ease with which it oxidises
to a ferric salt, is often used as a reducing agent. It is also
extensively used in the manufacture of black dyes, and of ink.

The commonest ferric salt is ferric cUoride, FeCls, which

246 INORGANIC CHEMISTRY

can be prepared in the anhydrous state by passing chlorine
over heated metallic iron. At a high temperature it sublimes
in the form of black scales. When this salt is dissolved in
water it forms a yellow solution which, on evaporation, de-
posits the ordinary yellow hydrate FeCl3,6H20. The solu-
tion of ferric chloride from which this hydrate may be obtained
is usually prepared by dissolving iron in hydrochloric acid, and
then oxidising the ferrous chloride thus formed by boiUng
with a little nitric acid. The equation for this action is —

3FeCl2 + 3HCI + HNO3 = 3FeCl3 + NO + 2H2O

Ferrous salts are almost invariably oxidised to ferric salts in
this way by means of nitric acid.

When sodium hydroxide is added to a solution of ferrous
sulphate quite free from ferric salt, a white precipitate of
ferrous hydroxide, Fe(0H)2, is produced —

FeS04 + 2NaOH == Fe(0H)2 + Na2S04

The precipitate obtained from an ordinary ferrous solution is
always dark green in colour, owing to the presence of a little
ferric salt in the ferrous solution.

Sodium hydroxide, when added to a solution of a ferric
salt, gives a reddish brown precipitate of ferric hydroxide —

FeCla + sNaOH = Fe(0H)3 + sNaCl

This hydroxide is the base corresponding to the ferric salts,
and when heated yields the basic oxide Fe203.

Between ferrous oxide, FeO, and ferric oxide, FcaOa, there
exists a ferroso-ferric oxide, Fe304, which is distinguished by
being magnetic like metallic iron. It occurs in nature, and
on account of its magnetic properties is sometimes known as
iodestone. It is not a basic oxide in the sense of having a
definite set of salts corresponding to it. When dissolved in
acids it yields a mixture of ferrous and ferric salts —

FcgOA + 4H2SO4 = FeS04 + Fe2(S04)3 + 4H2O

Ferrous sulphate Ferric sulphate

Ferric hydroxide is by no means a strong base, and solu-
tions of its salts are always acid on account of partial hydrolysis.
Like aluminium, it can neither form a carbonate nor a sul-

IRON 247

phide in aqueous solution. Thus, when a solution of sodium
carbonate is added to a solution of ferric chloride, ferric
hydroxide is precipitated, and carbon dioxide is evolved.

Fe2(S04)3 + sNagCOa + 3H2O =
2Fe(0H)3 + sNaaSO^ + 3CO2

Ferrous hydroxide is a considerably stronger base than
ferric hydroxide, and is capable of forming a carbonate. Thus,
when sodium carbonate is added to a pure solution of ferrous
sulphate ferrous carbonate is precipitated.

FeS04 + NagCOa = FeCOg + Na2S04
This carbonate occurs in nature as spathic iron.

CHAPTER XXXIII
CALCIUM— BABIUM

CALCIUM

Calcium occurs in nature chiefly as carbonate CaCOg, and
to a smaller extent as sulphate CaS04 and phosphate
Ca8(P04)2. An account of calcium carbonate has been
given in the introductory chapter, and it need only be
repeated here that calcium carbonate is chemically valuable
on account of the ease with which it can be converted into
the powerful basic calcium oxide CaO, which unites with
water to give the powerful though not very soluble base
calcium hydroxide Ca(0H)2.

Calcium sulphate occurs in nature chiefly as the hydrate
CaS04,2H20. Like the carbonate, this substance has various
forms. When opaque it is known ay gypsum and as alabaster^
when transparent and distinctly crystalline like calc-spar it is
known as selenite. When gypsum is heated to a temperature
somewhat below 200° it parts with three-fourths of its water
of crystallisation, and is converted into the lower hydrate
2CaS04,H2 0. This hydrate has the property of taking up
water again at the ordinary temperature to form the original
gypsum. It is therefore much employed under the name of
plaster of Paris for making plaster casts. It is ground to a
fine powder, and then mixed with water to a stiff" paste, which
may be forced into moulds while it is still soft. It then
gradually hardens and sets to a mass of what is practically
gypsum, all the details of the mould being reproduced.

Calcium phosphate^ Ca3(F04)2^ is chiefly useful as a source
of phosphoric acid and phosphorus (Chapter XXVIII), and
not on account of the calcium which it contains.

Calcium chloride is the commonest soluble salt of calcium.
It is obtained as a bye-product in many chemical manufactures,
the calcium which it contains being derived originally from

248

CALCIUM— BARIUM 249

limestone or chalk, and th^ chlorine which it contains being
derived originally from sodium chloride. It may be readily
prepared by dissolving calcium carbonate in hydrochloric acid,
and crystallising the solution. The crystals which separate
have the formula CaCl2,6H20. These, when heated to a
high temperature, lose their water of crystallisation, and yield
anhydrous calcium chloride CaCl2, which on account of the
ease with which it absorbs moisture is much used as an agent
for drying gases, and those liquids in which it is insoluble.
When exposed to the air both the ahhydrous calcium chloride
and the hydrated chloride absorb moisture, and ultimately
produce a solution of calcium chloride.

The oxide, hydroxide, and carbonates of calcium have
already been described in Chapters I and XXI. For
calcium carbide see p. 152.

BARIUM

Barium is an element which closely resembles calcium,
both in its mode of occurrence, and in its general chemical
properties. Its compounds are not, however, nearly so widely
distributed as those of calcium, and only occur in compara-
tively small quantities. It is found chiefly as the sulphate
heavy spar BaS04, ^^^ 21s the carbonate BaCOg.

Most barium compounds are juade from the natural sulphate.
As this substance is almost perfectly insoluble in water, it must
be converted into a soluble form before it can be transformed
into other barium compounds (compare p. 95). The barium
sulphate is reduced by means of carbon at a high temperature,
the equation being —

BaS04 + 4C = BaS + 4CO
The barium sulphide thus produced is soluble in water, and
can be converted into other barium compounds by means
of the appropriate acids.

Thus, if it is desired to prepare barium chloride, which is
the soluble salt of barium mostly in use, the sulphide or the
carbonate may be decomposed by hydrochloric acid, according
to the equations —

BaCOa + 2HCI = BaCla + H2O + CO2
BaS + 2HCI = BaClg + HgS
The barium chloride thereby produced may be crystallised out

250 INORGANIC CHEMISTRY

of the solution on evaporation in the form of the dihydrate
BaCi2,2H20, which is not nearly so soluble as calcium
chloride, and is not deliquescent.

Barium hydroxide, Ba(0H)2, may be prepared on the large
scale by passing carbon dioxide and superheated steam over
barium sulphide. Barium carbonate is probably first formed
according to the equation —

BaS + H2O + CO2 = BaCOa + HgS

and then decomposed by the superheated steam as follows : —

BaCO^ + H2O = Ba(0H)2 + CO2

the carbon dioxide being swept off in the current of steam.
Barium hydroxide is more soluble in water than calcium
hydroxide, and the saturated solution called baryta water is
often used instead of lime-water in testing for, and in estimat-
ing, the amount of carbon dioxide in air. When it is brought
into contact with carbon dioxide, barium carbonate is im-
mediately formed, according to the equation —

Ba(0H)2 + CO2 = BaCOg + H2O

and separates out as a precipitate.

Both barium hydroxide and barium carbonate resist the
action of heat more strongly than the corresponding calcium
compounds — that is, they may be heated to redness without
decomposition into barium oxide and water or carbon
dioxide. If it is desired to prepare barium oxide, this is best
done by heating the carbonate, not alone, but with carbon,
when the following action takes place : —

BaCOa + C = BaO + 2CO

The carbon of the carbonate is here got rid of, not as carbon
dioxide, as is the case with calcium, but as carbon monoxide.
Barium oxide, like calcium oxide, combines very readily with
water to form barium hydroxide. When heated in air it
absorbs oxygen, with formation of barium dioxide BaO 2* Its
use in preparing oxygen from the air has already been re-
ferred to (p. 128).

CALCIUM— BARIUM 25 1

If we consider in what points calcium and barium resemble
each other in their compounds, we find first of all that the
compounds have similar formulae. Not only, however, is there
this formal resemblance, but also a real resemblance in
chemical properties. Thus the carbonates and the sulphates
are nearly insoluble in water. The oxides unite very readily
with carbon dioxide, and with water ; but as has already been
mentioned, the resulting carbonate and hydroxide are much
less easily decomposed in the case of barium than in the case
of calcium. The hydroxides of both calcium and barium are
sparingly soluble in water, and the solutions which they form
are strongly alkaline. To distinguish them from the metals
of the alkalies which give freely soluble hydroxides, these
elements, together with strontium, which is intermediate in
properties between the two, are generally called metals of the
alkaline earths.

CHAPTER XXXIV
SODIUM— POTASSIUM— AMMONIUM

SODIUM

Practically all the sodium compounds are made from sodium
chloride^ which occurs abundantly in sea water, in brine springs,
and in some places as solid rock salt. It has already been
indicated that by electrolysing a solution of sodium chloride,
both chlorine and sodium hydroxide may be produced. This
electrolytic process is coming into extensive use as a source of
sodium hydroxide, and of the sodium carbonates which are
derived from sodium hydroxide by treatment with carbon
dioxide, according to the equations —

NaOH + CO2 = NaHCOa

Sodium hydrogen carbonate

2NaOH + CO2 = NaaCOs + HgO

Sodium carbonate

There are two other methods still in use, however, for obtaining
carbonates of sodium from sodium chloride. The oldest of
these is called the Le Blanc process, and the chemical actions
involved are the following: — First, the sodium chloride is
converted into sodium sulphate by the action of sulphuric
acid, which, as we have already seen, occurs in the following
two stages, the acid sulphate being first produced : —

NaCl + H2SO4 = NaHS04 + HCl

Sodium hydrogen sulphate

NaCl + NaHS04 = Na2S04 + -HCl

Sodium sulphate

The normal sodium sulphate thus obtained is reduced by
means of carbon at a high temperature with production of
sodium sulphide.

Na2S04 + 2C = Na2S + 2CO2

252

SODIUM— POTASSIUM— AMMONIUM 253

Limestone is added to the mixture of sulphate and carbon,
so that at the same time we have the reaction —

NaaS + CaCOa = NagCOg + CaS

When the product of the reaction is treated with water, im-
pure sodium carbonate is dissolved away from the calcium
sulphide, and is then subjected to various processes of puri-
fication. The still somewhat impure sodium carbonate which
has not been purified by crystallising, but merely by roasting,
is called soda ash. When the sodium carbonate is recrystal-
lised, it separates out as the decahydrate NagCOsjioHgO,
which is familiarly known as washing soda.

Another process, by which sodium hydrogen carbonate is pre-
pared, is also much employed. This is called the ammonia
soda process, and is based on the following action : — When
ammonium hydrogen carbonate and sodium chloride are
brought together in concentrated solution, the following de-
composition may take place : —

(NH4)HC03 + NaCl = NaHCO^ -h NH^Cl

In dilute solution all these salts would remain dissolved, but
when very little water is present the least soluble of the four —
namely, sodium hydrogen carbonate — falls out. In practice
this action is brought about by taking strong brine, saturating
it first with ammonia gas, and then leading carbon dioxide
through it until the sodium hydrogen carbonate falls out.

Sodium hydrogen carbonate, NaHCOs usually called bi-
carbonate of soda or baking soda, decomposes when heated at
a comparatively low temperature, according to the equation —

2NaHC03 = NagCOg + HgO -I- CO2

so that the normal carbonate^ Nag CO 3, can easily be prepared
from it. Normal sodium carbonate does not lose carbon
dioxide even at a red heat.

Practically all sodium compounds are made either during
the manufacture of the hydroxide and carbonates, or are
prepared by the action of acids on the hydroxide or
carbonates.

Thus sodimn snlpbate^ Na2S04^ is produced in the Le
Blanc process in the commercial form known as salt cake.

254 INORGANIC CHEMISTRY

When this is dissolved in water and the solution crystal-
lised the decahydrate Na2S04,ioH20 separates out. This
hydrate is known as Glauber's salt.

Sodium hydroxide or caustic soda, NaOH, can be made from
the carbonate by treating its solution with calcium hydroxide,
when double decomposition occurs, according to the equation —

NagCOs + Ca(0H)2 = CaCOg + 2NaOH

The soluble sodium hydroxide is separated from the insoluble
calcium carbonate and the solution evaporated. In order
that this reaction may take place, it is necessary that the
solution should not be too concentrated, otherwise the re-
action proceeds to a certain extent in the reverse direction —
namely,

CaCOa + 2NaOH = NagCOa + Ca(0H)2

On driving off the water from a solution of caustic soda, the
sodium hydroxide does not separate out in the crystalline state,
but fuses as the temperature rises. The fused caustic soda is
usually cast into sticks for laboratory purposes. Sodium
hydroxide readily absorbs both moisture and carbon dioxide
from the air, is excessively soluble in water, and very strongly
alkaline. It is not decomposed by heat at any temperature.

Metallic sodium is prepared from sodium hydroxide either
by electrolysis of the fused hydroxide, when the sodium travels
to the kathode and is there separated as the fused metal, or
by reduction at a high temperature by the carbon of iron
carbide.

6NaOH + 2C = 2Na2C03 + 3H2 + 2Na

The metallic sodium is in the state of vapour, so that it comes
off with the hydrogen and is condensed in iron tubes to a
liquid, which afterwards solidifies.

Sodium, though a metal, is lighter than water and so soft
that it can easily be cut with a knife. A freshly cut surface
has a bright, silver-white appearance, but it immediately
tarnishes on exposure to air, being attacked by the moisture
and the carbon dioxide of the air. When thrown into water,
even at the ordinary temperature, it at once attacks the water

SODIUM— POTASSIUM— AMMONIUM 255

with formation of sodium hydroxide and evolution of hydrogen,
according to the equation —

2Na + 2H2O = 2NaOH + Hg

The temperature of the reaction often rises so high that the
hydrogen takes fire and burns with a yellow flame, due to the
presence of sodium vapour. All sodium compounds when
heated to a high temperature impart this characteristic yellow
colour to a flame.

When sodium is heated in air it takes fire and burns with a
bright yellow flame, not to form, as we might expect, the basic
oxide NagO, but sodinm peroxide NagOg. As we have seen,
this sodium peroxide is really a salt of the feeble acid hydrogen
peroxide H2O2, which can be readily obtained from it by the
action of dilute acids (p. 168). Sodium peroxide is now pre-
pared on the commercial scale, and is useful as an oxidising
agent.

Sodinm nitrate^ NaNOs, otherwise known as Chili saltpetre^
occurs in quantity in the rainless districts of Chili and Peru.
It is used, not as a source of sodium, but as a source of
nitrates and nitric acid. It is much employed as a manure,
being a convenient source of nitrogen for plants.

POTASSIUM

Nearly all potassium compounds are now derived from the
salt deposits at Stassfurt, near Magdeburg, in Prussia. These
deposits have been formed by the evaporation of sea water
under peculiar conditions, the sodium chloride having ap-
parently crystallised out first, and the other salts, which occur
in smaller quantity in water, being then deposited separately.
These salts are chiefly magnesium salts and potassium salts.
Potassium chloride, as the mineral sylvine^ is present in com-
paratively small quantity, the chief source of potassium com-
pounds being a double chloride of potassium and magnesium
called carnallite^ which has the composition KCl,MgCl2,6H20.
As potassium chloride is less soluble than magnesium chloride,
this salt by proper treatment with water may be made to yield
crystals of potassium chloride.

Just as all the compounds of sodium are derived from sodium
chloride, the corresponding compounds of potassium may be

256 INORGANIC CHEMISTRY

produced by similar processes from potassium chloride, with
thfi exception that potassium hydrogen carbonate cannot be
produced from potassium chloride by a process analogous to
ammonia soda process.

When plants are burned, the ash which they leave behind
contains a very large proportion of potassium carbonate
K2CO3. This 'substance does not exist as such in the plants,
but is derived from the decomposition of potassium salts of
complex organic acids, which are contained in them. These
salts are necessary for all vegetable life, so that potassium
salts must be present in the soil in which plants grow. On a
natural soil, where the plants decay in the same place as that
on which they develop, the potassium salts in the plants
return to the soil. Where, however, a succession of crops is
grown on the same soil and removed year after year, the land
gets poorer and poorer in potassium salts, which must then be
replaced by adding potassium in some form. A potassium
compound, which is much used as a manure to supply potash
to plants, is the Stassfurt mineral kainite K2S04,Mg2S04,
MgCl2,6H20.

Potassium carbonate, E2CO3, used to be derived almost
entirely from wood ashes, and the name, pot ashes^ being
applied to the crude carbonate in a certain form, is the origin
of the terms potash and potassium. Potassium carbonate is
now chiefly made by a process analogous to the Le Blanc
process for the manufacture of sodium carbonate. From it
potassium hydrogen carbonate KHCO3 can be prepared by
the action of carbon dioxide, and potassium hydroxide KOH
can be prepared by the action of calcium hydroxide.

Potassium nitrate, ENO3, which is chiefly used in the pro-
duction of gunpowder, is now mostly prepared by the double
decomposition of potassium chloride and sodium nitrate. The
sodium nitrate itself cannot be used in the manufacture of
gunpowder, because it is slightly hygroscopic, so that in gun-
powder it would attract moisture, and thus make the powder
damp. Potassium nitrate is not hygroscopic, and gunpowder
containing it remains quite dry when exposed to the air. The
double decomposition between sodium nitrate and potassium
chloride is carried out at a somewhat high temperature, the
result being that sodium chloride, which is in these circum-
stances the least soluble salt, first falls out and leaves excess of

SODIUM— POTASSIUM— AMMONIUM 257

potassium nitrate in the solution. On cooling this potassium
nitrate crystallises. Reference to the curves on p. 15 shows
that whilst at a high temperature potassium nitrate is much
more soluble than sodium chloride, at a low temperature this is
not the case.

The metal potassium can be formed by the electrolysis of
fused potassium hydroxide, just as sodium can be formed by
the electrolysis of fused sodium hydroxide. Like sodium, it
is a soft very light metal which immediately tarnishes in air and
attacks water with great vigour, according to the equation —

2K + 2H2O = 2KOH + H2

The temperature of the reaction is so high that the hydrogen
is inflamed, burning with the lavender flame characteristic
of all potassium compounds.

The corresponding salts of potassium and sodium resemble
each other in appearance and properties very closely, the
principal difference between them being in their solu-
bility. In nearly all cases potassium or sodium salts may
be used indiscriminately; although, as a matter of practice,
sodium salts are always preferred to potassium salts on account
of their much smaller cost. Occasionally, of course, it happens
that a slight difference in solubility, or attraction for moisture,
may render it necessary to employ a potassium salt instead of
a sodium salt. An instance of this has just been given in the
case of gunpowder.

AMMONIUM

The resemblance between potassium salts and ammonium
salts, except in their behaviour towards heat, is extremely
close, much closer, in fact, than the resemblance of potassium
salts to sodium salts. This resemblance exists in spite of the
fact that the potassium salts contain a metal, and the ammonium
salts contain no metal, but a compound radical or group NH4,
the constituents of which are gases in the free state.

It has already been stated that the source of ammonium
compounds is ammonia derived from the distillation of coal
for the purpose of producing coal gas. The ammonia gas is
absorbed in water, in which it is extremely soluble, forming the
ammoniacal liquor of the gas works, which, when distilled with
R

258 INORGANIC CHEMISTRY

lime gives off ammonia gas in a purer form, which can then
be reabsorbed by sulphuric acid, according to the equation —

2NH3 + H2SO4 = (NHO2SO4

The crude ammonium sulphate^ (NH4)2S04^ is heated to
destroy some tarry material with which it is mixed, and then
purified by crystallisation. It is extensively used as a manure
for supplying nitrogen to the soil, in a form which plants can
assimilate.

Ammonium carbonate is prepared by heating a mixture of
ammonium sulphate and calcium carbonate in the form of
chalk. When this mixture is heated, gases come off which
condense again to form a white solid substance, which,
however, is not pure normal ammonium carbonate (NH4)2C03,
but a mixture of ammonium hydrogen carbonate (NH4)HC03,
and a substance called ammonium carbamate (NH4)C02NH2.
This mixture behaves practically in aqueous solution as
ammonium carbonate, and is generally known by that name.

Ammonium chloride, NH4C1^ is prepared by absorbing
ammonia gas from the ammoniacal liquor of the gasworks in
hydrochloric acid instead of in sulphuric acid. These sub-
stances combine, according to the equation —

NH3 + HCl = NH4CI

The ammonium chloride, or sal ammoniac^ which is thus
obtained, is first heated gently to destroy tarry material, and
then purified by sublimation. When perfectly dry, ammonium
chloride on heating passes into ammonium chloride vapour.
When the ammonium chloride, however, is not absolutely dry,
it dissociates on vaporisation into atnmonia and hydrochloric
acid. These gases, at the high temperature necessary for the
sublimation, exist side by side without combining. When
the mixture is cooled, however, the gases recombine, with re-
production of ammonium chloride.

The radical ammonium^ NH4J has never been isolated, and
it is doubtful if it has any existence apart from the negative
radicals with which it is combined in salts.

Ammonium salts, like sodium and potassium salts, are all
soluble in water, the actual solubilities resembling those of
the corresponding potassium compounds rather than those of

SODIUM— POTASSIUM— AMMONIUM 259

the corresponding sodium compounds. When ammonium
salts are heated, however, they undergo decompositions unlike
any decomposition which can take place with similar potassium
or sodium salts. The essence of this decomposition is that
the ammonium radical splits up with production of ammonia
gas, when the acid radical is a non-oxidising radical; and
with production of nitrogen, or an oxide of nitrogen, when
the acid radical is an oxidising radical. Thus, if we heat
ammonium phosphate, we obtain ammonia gas and phosphoric
acid, which remains as a non-volatile residue, the equation
being —

(NHJ3PO, = 3NH3 + H3PO4

= 3NH3 + H2O + HPO3

The phosphate radical here has no oxidising power. When
we heat ammonium nitrite or ammonium nitrate, on the other
hand, which contain the oxidising nitrite and nitrate radicals,
we obtain no ammonia gas, but in the first case nitrogen, and
in the second case nitrogen monoxide, according to the
equations —

NH4NO2 = N2 + 2H2O

NH4NO3 = N2O + 2H2O

When the acid, as well as the ammonia, is volatile, both acid
and ammonia come off together, and condense on cooling to
form the original ammonium salt. An example of this has
just been given in the case of ammonium chloride.

INDEX

Absolute temperature, io8
Absorption coefficient, 1 1 1
Acetates, 68, 72
Acetic acid, 72
Acetylene, 152
Acid, 40, 62

anhydrides, 53, 54

salts, 64

Acidic oxides, 53

Acids, action on metals, 55-57, 79

basicity, 65

common, 40, 45-47

mutual displacement, 70

Air, 27, 122-126
Alabaster, 248
Alcohol, 151
Alkali metals, 48, 252
Alkalies, 40, 48

common, 47-49

Alkaline earths, 48, 248
Allotropic forms, 130, 143, 193, 207
Aludels, 226
Alum, 239
Alumina, 238
Aluminium, 98, 238

compounds, 238-24 1

Amalgam, 226
Ammonia, 48, 112, 155
Ammonium hydroxide, 48, 74

nitrate, 161, 259

nitrite, 154, 259

salts, 258-259

Amorphous substances, 16

Anhydrous substances, 139

Anions, 59, 83

Anode, 83

Aqua fortis, 40

Argon, 112, 122, 124, 155

Atmosphere, 122-126

Avogaaro's principle, 114, 133, 167

Baking soda, 253
261

Barium compounds, 249, 250

monoxide, 128, 250

peroxide, 128, 168, 250

Baryta water, 126, 250
Bases, 40, 48, 63
Basic oxides, 53

salts, 65, 174

Basicity of acids, 65
Bauxite, 238

Bisulphide of carbon, 203
Black lead, 143
Blast furnace, 243
Bleaching action, 177

powder, 171, 176

Blowpipe, 36

Blue vitriol, 18, 25, 224

Bone ash, 206

charcoal, 145

Boyle's law, 108
Brass, 223
Bromine, 182
Bronze, 223, 234
Bunsen burner, 33

Calamine, 236
Calcium carbide, 152

carbonates, 5, 135, 138, 148

compounds, 248-9

hydroxide, 6

hypochlorite, 176

oxide, 6

phosphate, 206, 214

silicate, 206

sulphide, 201

Calc-spar, 5
Calomel, 138, 228
Carbides, 152, 243
Carbon, 27, 143

dioxide, 4, 6, 50, 112, 146

disulphide, 203

monoxide, 50, 148

Carbonates, 53, 57, 73

262

INDEX

Carbonic acid, 72, 147
Carnallite, 255
Cast-iron, 243
Catalytic agent, 129
Cathode, 83
Caustic alkalies, 40
Cerussite, 233
Chalk, 5
Charcoal, 144
Chemical change, i

conditions for, 8-12

reversible, 11

equations, 6, 22, 104

substances, 5

tests, I

Chili saltpetre, 255

China clay, 238

Chlorates, 68, 176, 178

Chloride of lime = bleaching powder

Chlorides, 58, 68

Chlorine, 80, 170

monoxide, 174

peroxide, 175

Choke-damp, 148
Clay, 238
Coal, 145

gas, 149, 150, 156, 167

Coefficient of absorption, ill
Coke, 145

Collection of gases, 118
Combination, 7
Combustibles, 39
Combustion, 27-39
Compounds, 19
Copper, 222

compounds, 223-225

pyrites, 192, 222

Copperas, 245
Corundum, 238
Corrosive sublimate, 227
Crystallisation, 16-18

water of, 18, 138

Cupellation, 231

Dalton's law of partial pressures, 121
Davy lamp, 38
Decomposition, 7
Deli(^uescence, 139
Density of gases, 114, 167
Desilverisation of lead, 231

Diamond, 143
Diffiision of gases, 1 1 7- 1 1 8
Dissociation, 158, 210, 258
Double decomposition, 67-75, 9'"93

Efflorescence, 139

Electrical charges, 60, 76

Electro-chemical tables, 78, 80

Electrodes, 83

Electrolysis, 82-89

Element, 19

Emery, 238

Equations, 6, 22-26, 104- 107

Equivalence, 41, 89

Ethylene, 151

Euchlorine, 180

Expulsion of acids from salts, 71

Faraday's law, 89
Ferric oxide, 28, 246

salts, 102, 245

Ferrous compounds, 102, 245

Fire-clay, 238

Fire-damp, 150

Flame, 31-39

Fluorine, 189

Formulae, 21

molecular, 114

Fuels, 27

Galena, 230

Galvanised iron, 237

Gas carbon, 145

Gases, 108- 121

G^y Lussac's law of expansion, 108

— of reacting volumes, 113

German silver, 223
Glauber's salt, 254
Gold, 80

Gram molecular volume, 114
Graphite, 143
Gun-metal, 234
Gunpowder, 9, 194, 256
Gypsum, 248

Haematite, 242

Halogens, 189

Hardness of water, 135, 148

Heat of combustion, 29, 30, 35

Heavy spar, 249

INDEX

263

Henry's law, 1 1 1
Hydrates, 138
Hydriodic acid, 187, 199
Hydrobromic acid, 183
Hydrocarbons, 32, 150
Hydrochloric acid, 47, 112, 172
Hydrogen, 27, 78, 165

peroxide, 168

Hydrolysis, 141
Hydrosulphides, 201
Hydroxides, 40, 53, 63, 68
Hygroscopic substances, 139
Hypochlorites, 175
Hypophosphites, 215
Hyposulphate of soda, 17, 200

Iceland spar, 5

Ignition point, 10, 37

Indicators — e.g. litmus, 40, 66

Insoluble substances, 68, 138

Iodic acid, 187

" Iodide of starch," 132, 187

Iodides, 188

Iodine, 185

lonisation, 76

Iron, 28, 242

pyrites, 191, 242

salts, 245

Kainite, 256
Kaolin, 238
Kathode, 83
Kations, 59, 84
Kipp's machine, 147

Lamp-black, 143
Lead, 230

compounds, 232-3

Lime, 1-7

water, 16, 28

Liquefaction of gases, 130
Litharge, 29, 232
Litmus, 40
Lodestone, 242
Luminosity of flame, 32
Lunar caustic, 221

Magnesium, 155
Magnetic iron ore, 242, 246
Malachite, 225

Manganese dioxide, 129, 170
Marble, 5
Marsh gas, 150
Matches, 207

safety, 208

Mercuric oxide, 52, 228

salts, 103, 226

Mercurous salts, 103, 227
Mercury, 52, 225
Metallic radicals, 59, 64
Metals, 54, 220

action on acids, 55-57, 79

water, 79, 165

occurrence in nature, 80

Metaphosphoric acid, 209, 212
Methane, 150
Milk of sulphur, 193
Mixed salts, 64
Mixture of gases, 117
Molecular formulae, 114

weights of gases, 114

Mond gas, 150
Mosaic gold, 235
Muriatic acid, 172

Negative radicals, 59

Neutralisation, 40-44, 67

Nickel, 223

Nitrates, 53, 57

Nitre, 14, 256

Nitric acid, 46, 99, 104, 157-164

anhydride, 157

oxide, 160, 196

Nitrites, 163
Nitrogen, 27, 122, 154

pentoxide, 157

peroxide, 157, 196

trichloride, 177, 181

trioxide, 159

Nitrous acid, 163

oxide, 161, 163

Nomenclature of salts and acids, 63
Non -metallic elements, 54, 220
Nordhausen acid, 197
Normal salts, 65

temperature and pressure, no

Notation, chemical, 19-20

Oil of vitriol, 45
Olefiant gas, 151

264

INDEX

Orthophosphoric acid, 214
Oxidation, 50-96

stage of, 51, loi

Oxides, 5054
Oxidising agents, 99-100
Oxygen, 27, 122, 126-133
Ozone, 130

Partial volumes and pressures, 120-

121
Perchlorates, 179
Perchromic acid, 169
Pewter, 234
Phosphates, 214
Phosphides, 208, 216
Phosphine, 216
Phosphonium iodide, 217
Phosphoric acids, 21 1-2 16

anhydride, 209

Phosphorous acid, 184, 215
Phosphorus, 28, 124, 206

bromides, 211

chlorides, 210

iodides, 211

oxides, 209

oxychloride, 211

Phosphuretted hydrogen, 216
Pipe-clay, 238
Plaster of Paris, 248
Plumbago, 143
Positive radicals, 59
Pot-ashes, 256
Potassium, 257

bichromate, 169

chlorate, 129, 179

hydroxide, 47

iodide, 100

nitrate, 129, 256

perchlorate, 180

permanganate, loi

salts, 255-257

Pressure, normal, ITO
Producer gas, 149
Pseudo-solution, 142
Pyrites, 191 ^
Pyrophosphoric acid, 212

Quicklime, 1-7
Quicksilver, 225

Radicals, 59-66
Reacting volumes, 112
Red lead, 129, 232
Reducing agents, 98-100
Reduction, 96
Respiration, 29, 130
Reverberatory furnace, 230
Reversible actions, 1 1
Rock salt, 252
Ruby, 238

Safety lamp, 38
Sal ammoniac, 258
Salt, 82, 252

cake, 253

decomposition, 57

formation, 42, 55

radicals, 59, 64

Saltpetre, 14, 256
Salts, 42, 63-65
Sapphire, 238
Saturated solutions, 13
Selenite, 248
Silica, 206
Silver, 220, 231

acetylide, 153

oxide, 169

salts, 221

Slag, 243
Slaked lime, 2-9
Soap-test, 135
Soda-ash, 253

caustic, 47, 254

Sodium, 79, 254

compounds, 252-255

hydroxide, 47, 254

iodate, 186

peroxide, 168

phosphates, 214

thiosulphate, 17, 200

Solder, 234
Solubility diagram, 15

of gases, III, 121, 123, 130

of salts, 68

Solutions, 9, 13-18
Spathic iron, 242
Stannic compounds, 234-235
Stannous compounds, 234-235
Steel, 244

INDEX

265

Strength of acids, 70-72
Strontium, 251
Sugar, 8

Sulphates, 58, 197
Sulphides, 73, 201
Sulphites, 195
Sulphur, 28, 191

chloride, 203

dioxide, 28, 51, 112, 194

trioxide, 51, 195

Sulphuretted hydrc^en, 72, 112, 200
Sulphuric acid, 45, 195, 198
Sulphurous acid, 194
Symbols, 19-20
Sympathetic ink, 140

Temperature of reaction, 10

absolute, 108

normal, 1 10

Test papers, 40
Tests, I

Thiosulphates, 199
Thiosulphuric acid, 200
Tin, 233

Tin compounds, 234-235
Tin-plate, 233

Unsaturated compounds, 152

Vapour density, 115
Vermilion, 228
Vitriol, oil of, 45, 198
Vitriols, 197
Volatility of acids, 7 1
Volumes, law of, 113

Washing soda, 253
Water, 134-142

vapour in air, 125

gas, 150

of crystallisation, 18, 138

Weights combining, 20

definite, 3

White lead, 233

precipitate, 228

Wrought iron, 244

Zinc, 236

compounds, 237, 238

TUK RIVBRKIUK PUtSS LIMITBD
BDINBURGH

Internet Archive Audio

Featured

Top

Images

Featured

Top

Software

Featured

Top

Books

Featured

Top

Video

Featured

Top

Mobile Apps

Browser Extensions

Archive-It Subscription

Save Page Now

Full text of "Elementary Inorganic Chemistry"

See other formats